;-NRLF 


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Db3 


IRATORY 


1INGED  TO 


CHE 


MCPH-!  FtSo 


THE   METRIC   SYSTEM 

1.  Length.    The  unit  is  the  meter.   It  is  equal  to  39.37  in, 
10  millimeters  (mm.)  =  1  centimeter  (cm.) 
10  centimeters  =  1  decimeter  (dm.) 

10  decimeters  =  1  meter  (m.) 

1000  meters  =  1  kilometer  (km.) 

The  only  measures  of  length  ordinarily  used  by  the 
chemist  are  the  millimeter  and  the  centimeter ;  thus,  the 
height  of  the  barometer  at  the  sea  level  is  recorded  at 
76  cm.  or,  more  commonly,  as  760  mm.,  and  not  7  dm. 
6  cm. 

~cn  2.  Volume.    The  unit  generally  used  is  the  cubic  centi- 

meter. 

1000  cubic  millimeters  =  1  cubic  centimeter  (cc.) 
1000  cubic  centimeters  =  1  cubic  decimeter  =  1  liter  (1.) 
1000  cubic  decimeters  =  1  cubic  meter 

The  chemist  uses  only  the  cubic  centimeter  and  the 
liter  as  measures  of  volume.  Thus,  the  volume  of  a  test 
tube  is  given  as  (say)  25  cc. ;  that  of  a  flask  as  (say) 
500  cc.  or  0.5  liter. 

3.  Weight.  The  unit  is  the  gram.  This  is  the  one  thou- 
sandth part  of  the  weight  of  a  piece  of  platinum  preserved 
at  Sevres,  France.  It  is  equal  to  15.43  grains. 

10  milligrams  (mg.)  =  1  centigram  (eg.) 
=       o  10  centigrams  =  1  decigram  (dg.) 

10  decigrams  =  1  gram  (g.) 

1000  grams  =  1  kilogram  (kg.) 

The  gram  is  practically  the  only  unit  used  by  the 
chemist.  Thus,  the  weight  of  a  crucible  is  given  as  (say) 
10.532  g.  and  not  10,532  mg.  or  10  g.  5  dg.  3  eg.  2  mg. 

Note  that 

1  pound  troy  =  5760  grains  =  373.24  grams 

1  pound  avoirdupois  =  7000  grains  =  453.59  grams 

1  U.S.  quart  =  0.95  liter 

Also  note  that 

1  centimeter  =  nearly  |  inch 

1  kilogram  =  nearly  2£  pounds  (avoirdupois) 


LABORATORY  MANUAL 

ARRANGED  TO  ACCOMPANY  THE 
SECOND  EDITION  OF 

"A  COURSE  IN  GENERAL  CHEMISTRY'7 


BY 

WILLIAM  McPHERSON 

AND 

WILLIAM  EDWARDS  HENDERSON 

PROFESSORS  OF  CHEMISTRY,  OHIO  STATE   UNIVERSITY 


GINN  AND  COMPANY 

BOSTON    •    NEW  YORK    •    CHICAGO    •    LONDON 
ATLANTA    •    DALLAS    •    COLUMBUS    •    SAN  FRANCISCO 


COPYRIGHT,  1915,  1921,  BY 
WILLIAM  MCPHERSON  AND  WILLIAM  EDWARDS  HENDERSON 


ALL  RIGHTS  RESERVED 

522.12 


gfte   gtfrenaum 

GINN  AND  COMPANY  •  PRO- 
PRIETORS •  BOSTON  •  U.S.A. 


PREFACE 

From  an  educational  point  of  view,  chemistry  is  really  the 
oldest  of  the  experimental  sciences.  The  problem  as  to  what 
laboratory  work  should  constitute  the  first  year's  course  is 
therefore  one  to  which  a  great  deal  of  thought  has  been 
devoted,  and  many  educational  experiments  have  been  made  in 
an  endeavor  to  solve  it.  Much  ingenuity  has  been  exercised 
in  the  development  of  suitable  laboratory  experiments  and  in 
the  invention  of  simple  apparatus  adapted  to  a  beginner's 
inexperience,  and  a  great  wealth  of  admirable  illustrative 
experiments  is  now  at  the  command  of  every  teacher. 

For  one  who  sets  about  the  task  of  arranging  an  experi- 
mental course  for  the  beginner,  there  remains  little  oppor- 
tunity for  originality  or  invention.  His  problem  is  rather 
one  of  selection.  Accordingly  this  laboratory  manual  lays  no 
claim  to  originality,  either  in  method  or  in  content.  It  has 
been  slowly  developed  in  connection  with  the  large  beginning 
classes  in  which  the  authors  have  been  interested,  and  has 
been  revised  and  reprinted  privately  a  number  of  times.  In 
response  to  many  requests  it  has  again  undergone  a  thorough 
revision  and  has  been  arranged  to  accompany  the  text  by  the 
authors,  entitled,  "A  Course  in  General  Chemistry." 

In  common  with  nearly  all  teachers  of  chemistry,  the 
authors  have  had  to  deal  with  the  fact  that  the  first  course 
in  college  comprises  students  who  have  had  an  elementary 
course  in  the  high  school,  as  well  as  those  who  have  had  no 
earlier  introduction  to  chemistry.  As  far  as  the  laboratory  is 
concerned,  the  authors  have  found  that  the  most  practical 
solution  of  this  problem  is  to  develop  a  manual  ample 
enough  to  meet  the  needs  of  both  classes  of  students.  The 

[ill]  / 


more  experienced  student  can  then  omit  the  exercises  with 
which  he  is  familiar,  and  the  student  with  less  experience 
can  omit  some  of  the  quantitative  exercises. 

Every  well-ordered  laboratory  has  its  own  system  of  sup- 
plies, stock  solutions,  desk  reagents,  and  locker  equipment, 
and  the  directions  in  a  manual  will  not  always  harmonize 
with  this  system.  For  example,  in  many  cases  in  which  the 
student  is  directed  to  prepare  solutions  or  mixtures  for 
specific  purposes,  it  may  be  much  better  to  have  a  supply 
on  the  side  shelf,  properly  labeled,  for  general  use.  In  a 
number  of  experiments,  particularly  those  involving  appa- 
ratus that  is  not  a  part  of  the  locker  equipment,  two  or  even 
more  students  may  work  together  to  advantage.  The  time 
at  the  disposal  of  the  class  will  not  always  permit  each 
student  to  complete  all  of  the  exercises,  and  variety  is  added 
by  assigning  parallel  experiments  to  alternate  students.  All 
such  adjustments  are  left  to  the  instructor. 

In  the  Appendix  will  be  found  a  number  of  suggestions 
to  the  instructor,  relating  to  the  details  of  special  apparatus 
or  reagents.  Suggestions  are  also  offered  as  to  a  suitable 
locker  equipment  and  as  to  the  apparatus  that  the  student 
will  need  from  tune  to  time.  Articles  may  obviously  be 
transferred  from  one  list  to  another  according  to  the  funds 
available  or  the  capacity  of  the  lockers.  A  list  of  the 
reagents  called  for  in  the  manual  is  provided,  together  with 
an  estimate  of  the  quantities  required  for  a  class  of  ten 
students. 

A  large  page  has  been  chosen  for  the  manual,  since  it 
makes  a  convenient  book  for  filing  in  laboratory  pigeonholes 
and  for  handling  during  correction.  The  blank  pages  can 
be  used  for  a  full  record  of  laboratory  observations,  or  more 
elaborate  reports  may  be  written  from  the  recorded  notes  if 
the  instructor  so  desires. 

OHIO  STATE  UNIVERSITY 
COLUMBUS,  OHIO 


PREFACE  TO  SECOND  EDITION 

The  publication  of  a  second  edition  of  the  authors'  "  A 
Course  in  General  Chemistry "  has  made  it  necessary  like- 
wise to  revise  the  Laboratory  Manual  arranged  to  accompany 
the  text.  Attention  is  called  especially  to  the  fact  that  it 
is  not  expected  that  any  one  student  will  perform  all  the 
experiments  included  in  the  manual.  Each  instructor  will 
select  such  as  best  meet  the  needs  of  his  class. 

OHIO  STATE  UNIVERSITY 
COLUMBUS,  OHIO 


CONTENTS 

CHAPTER  PAGE 

I.    MANIPULATION  AND  FUNDAMENTAL  PRINCIPLES  .     .  1 

II.    OXYGEN 11 

III.  HYDROGEN 15 

IV.  WATER  AND  HYDROGEN  PEROXIDE 24 

V.    THE  STATES  OF  MATTER 30 

VI.    THE  LAWS  OF  CHEMICAL  COMBINATION 32 

VII.    CARBON  AND  CARBON  DIOXIDE 36 

VIII.    NITROGEN  AND  THE  ATMOSPHERE 40 

IX.    SOLUTIONS 43 

X.    CHLORINE;    HYDROGEN   CHLORIDE;    HYDROCHLORIC 

ACID 46 

XI.    SODIUM  ;  SODIUM  HYDROXIDE                i 50 

XII.    ACIDS;  BASES;    SALTS;    NEUTRALIZATION    ....  51 

XIII.  IONIZATION 53 

XIV.  COMPOUNDS  OF  NITROGEN 57 

XV.    EQUILIBRIUM 65 

XVI.    SULFUR  AND  ITS  COMPOUNDS 68 

XVII.    THE  CHLORINE  FAMILY 75 

XVIII.   SOME  COMPOUNDS  OF  CARBON 80 

XIX.    THE  LAW  OF  GAY-LUSSAC 84 

XX.    COMBINING  WEIGHTS  AND  MOLECULAR  WEIGHTS     .  88 

XXI.    SOME  HYDROCARBONS 91 

XXII.    FUEL  GASES;  FLAMES;  MEASUREMENT  OF  HEAT      .  93 

XXIII.  CARBOHYDRATES;  ALCOHOLS;  SOAPS 97 

XXIV.  PHOSPHORUS  FAMILY 100 

XXV.   BORON  AND  SILICON 106 

XXVI.    COLLOIDS 108 

XXVII.    GENERAL    METHODS    FOR    PREPARATION    OF     COM- 
POUNDS    112 

XXVIII.    THE  ALKALI  METALS 114 

XXIX.    THE  ALKALINE  EARTH  METALS 118 

XXX.    THE  MAGNESIUM  FAMILY 120 

XXXI.    ALUMINIUM 123 

XXXII.    THE  IRON  FAMILY 126 

XXXIII.  COPPER;  MERCURY;  SILVER 130 

XXXIV.  TIN  AND  LEAD 134 

XXXV.    MANGANESE  AND  CHROMIUM 138 

APPENDIX 141 

TABLE  OF  ATOMIC  WEIGHTS Inside  Back  Cover 

[vi] 


LABORATORY  MANUAL 

ARRANGED  TO  ACCOMPANY 
"A  COURSE  IN  GENERAL  CHEMISTRY" 


LABOKATOKY  MANUAL 

CHAPTER  I 
MANIPULATION  AND  FUNDAMENTAL  PRINCIPLES 

1.  The  Bunsen  burner,  a.  The  Bunsen  burner  is  the  common 
form  of  apparatus  used  in  laboratory  operations  where  heat  is 
applied.  It  consists  of  the  tube  A  (Fig.  1),  screwed  into  the 
base  C.  The  tube  has  two  small  round  holes  near  its  lower 
part.  A  band  B,  provided  with  similar  holes,  fits  around  the 
lower  part  of  the  tube  in  such  a  way  that  the  holes  in  the 
tube  may  be  closed  or  opened  by  turning 
the  band.  Gas  is  admitted  to  the  burner 
through  D  by  means  of  rubber  tubing. 
Unscrew  the  tube  and  examine  the  differ- 
ent parts  of  the  burner;  then  put  them 
together  again  and  light  the  gas  by  holding 
a  lighted  match  4  or  5  cm.  above  the  tube 
and  turning  on  the  gas.  The  gas  should 
be  adjusted  so  as  to  give  a  flame  about 
10  cm.  high.  The  gas  entering  the  burner  FIG.  1 

mixes  with  air  drawn  in  through  the  holes 
in   the  lower   part  of    the    tube  and  burns  with  an  almost 
nonluminous  flame.    If  the  band  is  adjusted  so  as  to  close  the 
openings,  the  flame  becomes  luminous.    Always  use  the  non- 
luminous  flame  unless  otherwise  directed. 

b.  Hold  an  iron  wire  in  various  parts  of  the  flame  to  gain 
an  idea  as  to  the  relative  temperatures.  Pass  a  piece  of 
glazed  paper  transversely  through  the  flame,  holding  it  steady 
for  an  instant  across  the  flame.  What  information  do  you 


get  from  the  way  in  which  the  paper  is  scorched?  Pass  a 
piece  of  paper  edgewise  through  the  axis  of  the  flame  in 
the  same  way. 

c.  From  these  experiments  draw  a  diagram  of  the  flame, 
indicating  the  hot  and  the  cooler  regions. 

2.  Heating  a  liquid  in  a  test  tube.  Hold  the  test  tube 
between  the  thumb  and  fingers  (Fig.  2),  constantly  rotating 
it  backward  and  forward  so  as  to  apply  the  heat  uniformly. 
The  heat  should  be  applied  to  the  upper  portion  of  the  liquid, 
care  being  taken,  however,  that  the  flame  does  not  strike  the 
tube  above  the  level  of  the  liquid.  In  case  the  upper  part  of 


FIG.  2 


FIG.  3 


FIG.  4 


the  tube  becomes  heated  it  may  be  supported  by  a  test-tube 
holder  (Fig.  3)  or  by  a  band  of  paper  circling  the  upper  part 
of  the  tube  (Fig.  4). 

The  sudden  formation  of  vapor  at  the  bottom  of  the  tube 
sometimes  causes  the  contents  of  the  tube  to  be  thrown  out ; 
hence  care  must  be  taken  not  to  point  the  tube  toward  anyone. 

Half  fill  a  test  tube  with  water  and  apply  heat  until  the 
water  boils  rapidly. 

3.  Pouring  a  liquid  from  one  vessel  to  another.  In  this 
operation  care  must  be  taken  to  prevent  the  liquid  from  run- 
ning down  the  side  of  the  vessel  from  which  it  is  poured. 
A  glass  rod  should  be  held  lightly  against  the  rim  of  the 
vessel,  as  shown  in  Fig.  7.  The  liquid  will  flow  down  the  rod. 
Fill  a  beaker  with  water  and  transfer  it  slowly  to  another 

[2] 


vessel  without  using  the  glass  rod;  repeat,  using  the  glass 
rod.  What  difference  is  noted  ? 

In  pouring  a  liquid  from  a  bottle  a  glass  rod  may  be  used ; 
or  the  neck  of  the  bottle  may  be  placed  lightly  against  the  rim 
of  the  vessel  into  which  the  liquid  is  being  poured  (Fig.  5). 
This  will  prevent  the  liquid  from  running  down  the  side  of 
the  bottle.  Obviously  the  stopper  should  never  be  laid  down 
on  the  desk.  It  should  be  caught  between  the  fingers,  as 
shown  in  Fig.  6.  This  leaves  the  hand  free  to  grasp  the  bottle, 
as  shown  in  Fig.  5. 

4.  Decantation.  It  is  often  necessary  to  separate  a  liquid 
from  a  finely  divided  solid  which  is  suspended  in  it.  This 


FIG.  5 


FIG.  6 


may  be  done  most  simply  by  decantation.  The  operation  con- 
sists in  allowing  the  solid  to  settle  and  then  pouring  off  the 
liquid.  The  method  can  be  used  only  when  the  solid  is  heavy 
and  readily  settles  in  the  liquid.  Shake  up  some  fine  sand  in 
water,  then  allow  it  to  settle  and  decant  the  water. 

5.  Filtration.  As  a  rule  the  solid  will  not  readily  settle  or 
will  do  so  only  after  long  standing.  In  such  cases  the  mixture 
is  filtered ;  that  is,  poured  on  a  filter  paper,  which  allows  the 
liquid  to  run  through  but  retains  the  solid.  To  prepare  the 
filter  paper,  fold  it  along  a  diameter  into  halves,  then,  at  right 
angles  to  the  first  fold,  into  quarters.  The  folded  filter  is  then 
opened  so  as  to  form  a  cone,  half  of  which  is  composed  of 
three  thicknesses  of  paper  and  the  remainder  of  one  thickness. 
Fit  the  cone  into  a  funnel,  of  such  a  size  that  the  paper  does 

[3] 


FIG.  7 


not  quite  reach  the  top.  The  paper  must  accurately  fit  the 
funnel ;  if  it  does  not,  make  it  do  so  by  varying  the  fold. 
Place  the  paper  in  the  funnel  and 
thoroughly  wet  it  with  water. 
After  the  water  has  run  through, 
press  the  paper  firmly  against  the 
sides  of  the  funnel  with  the  finger 
so  as  to  remove  any  air  bubbles  be- 
tween the  paper  and  the  glass.  The 
filter  is  now  ready  for  use. 

The  process  of  filtration  not  only 
enables  us  to  separate  liquids  from 
solids  but  also  certain  solids  from 
each  other.  To  illustrate  this  fact, 
grind  a  piece  of  chalk  to  a  powder 
in  a  mortar  and  mix  the  product 
intimately  with  about  an  equal  bulk  of  common  salt.  Now 
separate  the  two  solids  as  follows:  Place  the  mixture  in  a 
small  beaker,  add  about  50  cc.  of  distilled  water,  and  stir 
with  a  glass  rod.  (The  sharp  ends  of 
the  rod  must  be  rounded  by  rotating 
them  in  a  flame,  otherwise  the  beaker 
may  be  scratched  and  broken.)  The 
salt  dissolves,  forming  a  solution.  Filter 
off  the  insoluble  chalk,  collecting  the 
filtrate  (the  clear  liquid  which  passes 
through  the  filter  paper)  in  a  beaker 
(Fig.  7).  The  salt  may  be  recovered 
from  the  filtrate  by  the  process  of  evap- 
oration. To  perform  this  operation,  pour 
the  filtrate  into  an  evaporating-dish, 
then  support  the  dish  on  a  ring  stand 
(Fig.  8)  and  heat  gently.  The  tip  of 
the  flame  should  not  quite  touch  the  dish.  The  liquid  may  be 
made  to  simmer,  but  should  not  be  made  to  boil  violently 
(why?).  Withdraw  the  heat  as  soon  as  the  water  is  evapo- 
rated. Note  the  residue  left.  Convince  yourself  that  it  is  salt. 


FIG.  8 


6.  The  metric  system,    a.   Length.  By  means  of  the  scale 
on  the  front  cover  of  the  notebook  measure  the  length  (in 
centimeters)  of  various  pieces  of  apparatus  included  in  your 
outfit,  as  a  test  tube,  file,  and  blowpipe.    What  is  the  diam- 
eter of  your  filter  paper?    Estimate  the  lengths  of  various 
objects,   as  a  pencil,  a  test  tube;  then  measure.    Continue 
until  you  can  approximate  the  lengths  of  small  objects. 

b.  Volume.  By  means  of  a  graduated  test  tube  or  cylinder 
measure  (in  cubic  centimeters)  the  volumes  of  various  test 
tubes,  beakers,  and  flasks  included  in  your  outfit.  (In  reading 
the  volume  of  the  liquid  in  a  graduated  tube  always  read 
from  the  lower  part  of  the  meniscus  ;  that  is,  the  curved  surface 
of  the  liquid.) 

7.  The  balance.    The  weight  of  a  body  is  determined  on  a 
balance,  which  must  be  adapted  to  the  degree  of  precision 
required.     For   the   experiments  undertaken   in  this  course 
the  balance  must  weigh  accurately  to  one  one-hundredth  of 
a  gram,  and,  if  possible,  to  one  one-thousandth.    It  is  conven- 
ient to  remember  that  a  gram  is  approximately  one  thirtieth 
of  an  ounce. 

To  be  of  service  for  even  moderately  accurate  weighing,  a 
balance  must  be  used  with  the  greatest  care.  The  following 
directions  must  be  faithfully  observed : 

1.  At  all  times,  save  during  an  actual  weighing,  the  beam 
must  be  lifted  from  its  bearings  by  turning  the  thumbscrew 
at  the  base  of  the  case. 

2.  All  changing  of  objects  to  be  weighed  and  of  weights 
must  be  done  after  the  beam  has  been  lifted.    Never  take  off 
a  weight  while  the  beam  is  swinging. 

3.  Never  handle  the  weights  in  any  manner  save  with  the 
pliers  provided  with  the  weights.    Be  careful  not  to  dent  the 
weights  by  rough  use  with  the  pliers. 

4.  Keep  the  balance  pans,  the  case,  and  the  balance  shelf 
scrupulously  clean. 

5.  Never  weigh  chemicals  directly  on  the  pan  of  the  balance. 
First  weigh  a  suitable  container,  —  a  watch  glass  or  a  small 
beaker, —  then  place  the  material  in  this  and  again  weigh. 

[5] 


8.  The  process  of  weighing,    a.  See  that  upon  careful  lower- 
ing of  the  beam  the  vibrations  of  the  pointer  are  equal  on 
both  sides  of  the  zero  point.    If  there  is  any  considerable 
deviation,  have  the  instructor  adjust  the  balance. 

b.  Place  the  object  to  be  weighed  on  the  left  pan  and  the 
weights  you  judge  to  be  sufficient  on  the  right  pan.    Turn 
the  arrest  screw  very  slowly  until  you  see  in  which  direc- 
tion the  pointer  starts  to  move.    Arrest  the  beam  and  make 
a  new  adjustment  of  weights,  repeating  the  operation  until 
the  desired  equilibrium  point  is  found. 

c.  Add  up  the  weights  lying  on  the  pan,  and  verify  the 
count  by  the  empty  spaces  in  the  box.     Always  have  the 
weights  either  on  the  pan  or  in  the  box,  and  never  lay  them 
on  the  floor  of  the  balance  case. 

d.  Record  your  weighings  directly  in   your  notebook  — 
never  on  scraps  of  loose  paper.    Observe  that  the  weights 
are  in  decimals  and  that  their  sum  should  be  expressed  in 
grams  and  decimal  fractions  thereof. 

e.  Select  a  clean  watch  crystal  or  a  small  beaker  which  you 
will  be  using  for  weighing.    Carefully  weigh  it  and  record  the 
weight  in  your  notebook.    With  a  sharp  file  make  a  slight 
scratch  upon  the  beaker  just  below  the  lip.    Fill  the  beaker 
with  distilled  water  from  a  graduated  measure  and  note  the 
volume  added.    Weigh  the  beaker  and  water  and  determine 
the  weight  of  the  water.   How  does  the  weight  in  grams  com- 
pare with  the  volume  in  cubic  centimeters  ?    Why  is  this  ? 
Make   a  mark   on   a  test   tube   about   2  cm.   from  the  top. 
Weigh  the  empty  tube  and  also  the  tube  filled  with  water 
to  the   mark.    Determine  the  volume  to  the  mark.    Make 
use  of  the  tube  in  measuring  out  liquids  until  you  grow 
accustomed  to  estimating  volumes  in  cubic  centimeters. 

9.  Manipulation  of  glass  tubing,    a.   To   cut  glass  tubing. 
Place  the  tubing  on  the  desk  and  draw  the  edge  of  a  trian- 
gular file  across  the  point  at  which  you  wish  to  cut  the  glass. 

NOTE.  In  all  operations  requiring  the  application  of  a  strong  heat  to 
glass  the  heat  must  be  applied  gently  at  first ;  also,  the  highly  heated 
glass  must  be  cooled  slowly  to  anneal  it  properly. 

[6] 


After  the  glass  is  scratched,  take  the  tube  in  the  hands  with 
the  thumbs  placed  near  together,  just  back  of  the  scratch 
(Fig.  9),  and  gently  pull  the  glass  apart,  at  the  same  time 
exerting  a  slight  pressure  with  the  thumbs.  Do  not  try  to 
break  the  tubing  as  you  would  break  a  stick,  for  this  will 


FIG.  9 


FIG. 10 


splinter  the  glass  and  will  never  give  a  square  end.  If 
the  tube  does  not  yield  readily  to  the  gentle  pressure,  a 
deeper  scratch  must  be  made.  In  the  case  of  large  tubing 
it  may  be  found  necessary  to  file  a  groove  around  the  tube. 
The  edges  of  the  cut  tube  will  be  sharp.  They  should  be 
rounded  by  rotating  them  in  the  tip  of  the  Bunsen  flame. 
This  process  is  called  fire-polishing. 

b.  To  bend  glass  tubing.  Use  the  luminous  Bunsen  flame, 
spread  out  by  means  of  the  so-called  "  wing-top  "  burner 
(Fig.  10).  Hold  the  tube  lengthwise  in  the  flame,  gently 
rotating  it  so  that  all  sides  may  be  equally  heated.  Continue 
the  heating  until  the  glass 
easily  bends,  then  quickly  re- 
move it  from  the  flame  and 
bend  to  the  desired  shape 
(Fig.  11,  A).  Great  care  must 
be  taken  to  heat  the  tube  uni- 
formly, otherwise  the  bore  of 
the  tube  will  be  contracted 


FIG.  11 


(Fig.  11,  ^),  forming  a  bend 
which  not  only  is  unsightly 
but  is  easily  broken. 

c.   To  insert  a  glass  tube  in  a  cork.    The  cork  should  be  of 
such  a  size  that  the  smaller  end  will  just  enter  the  flask  or 

[7] 


FIG. 12 


bottle  in  which  it  is  to  be  used.  Soften  the  cork  by  rolling 
it  between  the  desk  and  a  piece  of  wood.  To  insert  a  glass 
tube  select  a  borer  slightly  smaller 
than  the  tube.  Place  the  cork  on 
the  desk  and  cut  about  halfway 
through  it,  not  by  punching,  but 
by  rotating  the  borer  under  gentle 
pressure  (Fig.  12).  Reverse  the  posi- 
tion of  the  cork  and  cut  through 
from  the  other  end.  Care  must  be 
taken  to  keep  the  borer  at  right 
angles  to  the  cork.  The  hole  should 

be  straight  and  smooth.  The  glass  tube,  rounded  at  the  edges, 
is  now  inserted  by  a  gentle  screwlike  motion.  If  the  hole  is 
too  small  to  admit  the  tube 
when  a  gentle  pressure  is 
applied,  it  may  be  slightly 

enlarged  with  a  round  file.     ?  — — • — ~ — • * 

The  tube  slips  in  better  if  FlG>  13 

both  cork  and  tube  are  wet. 

d.  To  draw  out  a  glass  tube  to  a  small  lore.    Heat  a  portion 
of  the  tube  in  the  Bunsen  flame  until  the  walls  of  the  heated 
portion  thicken  and  the  size  of  the 

bore  diminishes  (Fig.  13,  A).  The 
tube  must  be  constantly  rotated  to 
prevent  the  softened  portion  from 
sagging.  Now  quickly  remove  the 
tube  from  the  flame,  and,  holding  it 
in  a  vertical  position,  gently  pull  the 
ends  apart  until  the  bore  is  of  the 
desired  size  (Fig.  13,  B).  A  glass  jet 
may  be  formed  by  cutting  the  tube  at 
B,  and  rounding  the  edges  in  a  flame. 

e.  After     practicing     the     above 
manipulations,  make   a  wash  bottle 
according  to  Fig.  14.    Use  a  500-cc. 

flask.    A  and  B  are  glass  tubes.     B  is  connected  with  the 


FIG. 14 


glass  jet  D  by  a  piece  of  rubber  tubing  C.  The  edges  of 
the  glass  tubing  must  be  rounded.  When  the  wash  bottle 
has  been  approved  by  the  instructor,  fill  it  with  distilled 
water  and  set  it  aside  for  future  use. 

10.  Chemical  action,    a.  Hold  a  piece  of  iron  wire  in  the 
Bunsen  flame  for  a  few  seconds.    Is  the  iron  changed  ?    Exam- 
ine it  when  it  has  cooled.   Have  the  properties  changed  ?   Has 
a  chemical  action  occurred  ? 

b.  Repeat  a,  using  a  splint  of  wood  in  place  of  the  iron 
wire.    How  does  the  change  produced  differ  from  that  in  a? 
Has  a  chemical  action  occurred? 

c.  Place  enough  sugar  in  a  clean,  dry  test  tube  to  cover 
the  bottom  to  a  depth  of  1  cm.    Heat  it  gently  in  the  tip  of 
the  flame  as  long  as  any  changes  are  produced.    Note  all  the 
changes.    Is  the  product  sweet  ?    Is  it  soluble  in  water  ?    Do 
any  properties  remain  unchanged  ?    What  grounds  do  you 
have  for  assuming  that  a  chemical  action  has  taken  place  ? 

d.  Place  about  1  g.  of   common   salt  in   a  test  tube  and 
dissolve  it  in  a  little  water.     Pour  the  clear  solution  into 
an  evaporating-dish  and  evaporate  to  dryness.    What  is  the 
solid  left?    How  do  its  properties  correspond  to  those  of  the 
original  salt? 

e.  Cover  a  small  piece  of  zinc  in  a  test  tube  with  about 
5  cc.  of  water  and  add  carefully  3  or  4  drops  of  sulfuric 
acid.    Notice  that  the  zinc  dissolves  with  the  evolution  of  a 
gas.   Hold  a  burning  splint  at  the  mouth  of  the  test  tube  and 
note  the  result.   After  the  action  has  entirely  ceased,  filter  off 
any  undissolved  zinc  and  evaporate  the  solution  to  dryness 
(hood)  as  in  d.    How  does  the  change  differ  from  that  in  d  ? 
Distinguish  between  the  two  examples  of  solution. 

11.  Elements;  compounds;  mixtures,    a.  What  is  an  ele- 
ment ?    Are  iron  and  sulfur  included  in  the  list  of  elements  ? 
Weigh  out  separately  (on  paper)  2  g.  of  sulfur  and  2  g.  of 
clean  iron  filings  and  make  a  careful  list  of  their  properties. 
Try  the  effect  of  a  magnet  on  each.    Now  mix  the  two,  and 
grind  them  together  intimately  in  a  mortar.    Examine  the 
product  with  a  magnify ing-giass.     Can  you  distinguish  the 

[9] 


iron  from  the  sulfur  ?  Can  you  separate  them  with  a  mag- 
net ?  Have  they  undergone  any  change  in  properties  ?  What 
is  such  a  material  called  ? 

b.  Now  place  the  product  in  a  clean  test  tube  and  heat 
gently.  As  soon  as  the  mass  begins  to  glow,  quickly  withdraw 
the  tube  from  the  flame.  Does  the  mass  continue  to  glow  ? 
Now  heat  it  strongly  for  one  or  two  minutes ;  then  cool  the 
tube,  break  it,  and  examine  the  product  with  a  magnifying- 
glass.  Can  you  now  distinguish  between  the  iron  and  the 
sulfur  ?  Try  the  effect  of  the  magnet.  Of  what  is  the  sub- 
stance composed  ?  When  elements  combine  chemically,  do 
they  retain  their  original  properties?  What  is  the  product 
of  such  a  combination  called? 


CHAPTER  II 
OXYGEN 

12.  Collection  of  gases.   Fill  a  wide-mouthed  bottle  (250-cc.) 
with  water.    Cover  its  mouth  with  a  glass  plate,  being  careful 
to  exclude  all  air  bubbles.    Hold  the  plate  firmly  in  place, 
invert  the  bottle,  and  bring  its  mouth  below  the  surface  of 
the  water  contained  in  a  suitable  vessel.   The  most  convenient 
vessel  is  a  rectangular  box  of  sheet  iron  called  a  pneumatic 
trough  (Fig.  16).     Remove  the  glass  plate.     Why  does  the 
water  remain  in  the  bottle  ?    Now  fill  the  bottle  with  exhaled 
air  by  placing  one  end  of  a  piece  of  glass  or  rubber  tubing 
under  the  mouth  of  the  bottle  and  blowing  gently  through 
the  other  end.    The  bubbles  will  rise  and  fill  the  bottle. 

Before  the  bottle,  so  filled,  is  removed  from  the  trough, 
cover  its  mouth  tightly  with  a  glass  plate.  The  bottle  so 
covered  may  then  be  placed  on  the  desk  either  right  side  up 
or  in  an  inverted  position.  (When  should  it  be  placed  in 
an  inverted  position  ?) 

Fill  a  bottle  with  exhaled  air  and  transfer  the  air  so  collected 
to  another  bottle.  Draw  a  diagram  to  illustrate  a  suitable 
method  of  doing  this. 

13.  Preparation  of  oxygen,  a.  In  the  bottom  of  a  clean,  dry 
test  tube  place  about  0.5  g.  of  mercuric  oxide.    This  is  best 
accomplished  by  placing  the  oxide 

near  the  end  of  a  narrow   strip  of 

folded    paper    and    introducing    it 

carefully   into  the  tube,  as  shown  FIG.  15 

in  Fig.  15.    On  inclining  the   tube 

and  gently  tapping  the  paper,  the  oxide  will  be  deposited  in 

the  bottom  of  the  tube.  The  paper  is  then  withdrawn,  leaving 

' 


_ 


.   : 


the  sides  of  the  tube  perfectly  clean.  Hold  the  tube  between 
the  thumb  and  fingers  (Fig.  2)  and  apply  a  gentle  heat  to  the 
oxide.  The  tube  must  be  rotated  constantly  to  distribute  the 
heat;  otherwise  the  glass  may  soften.  While  the  heating  is 
maintained,  insert  a  glowing  splint  from  time  to  time  into 
the  mouth  of  the  tube.  Note  the  result.  Continue  to  heat  as 
long  as  any  gas  is  evolved.  What  remains  in  the  tube  ?  How 
has  the  heat  affected  the  mercuric  oxide  ? 

b.  In  a  similar  way  try  heating  other  oxides  such  as  iron 
oxide,  manganese  dioxide,  lead  peroxide,  barium  dioxide,  and 
copper  oxide.  Do  all  of  these  yield  oxygen  ? 

14.  Catalysis.    Place  about  1  g.  of  potassium  chlorate  in  a 
test  tube  and  carefully  heat  it  in  a  small  Bunsen  flame  until 
it  has  just  melted.    Can  you  see  any  gas  being  given  off  ? 
Continue  the  heating  until  you  can  plainly  see  that  gas  is 
liberated.    What  is  it  ?    From  the  time  of  heating  would  you 
suppose  that  the  temperature  had  been  raised  much  above 
the  melting  point?    Allow  the  liquid  to  cool  until  it  begins 
to  solidify,  and  then  just  melt  it  again.   Then  add  a  very  little 
powdered  manganese  dioxide,  dropping  it  in  from  a  small 
knife  blade.    What  effect  is  produced  ?    In  preparing  oxygen 
from  the  chlorate  what  would  be  the  advantage  of  adding 
manganese  dioxide  before  heating  ?    In  place  of  manganese 
dioxide  try  the  effect  of  iron  oxide  and  of  powdered  sand. 

15.  Usual  laboratory  method.  Arrange  an  apparatus  accord- 
ing to  Fig  16,  in  which  B  is  a  hard-glass  test  tube,  D  a  glass 
tube  bent  at  right  angles,  and  C  a  piece  of  rubber  tubing. 
Mix  intimately  on  paper  6  g.  of  potassium  chlorate  and  3  g. 
of  manganese  dioxide.    Transfer  the  mixture  to  a  hard-glass 
tube  and  insert  the  cork ;  then,  holding  the  burner  in  the 
hand,  heat  the  mixture  gently  with  a  small  flame,  applying 
the  heat  at  first  to  the  upper  part  of  the  mixture.    The  flame 
must  not  strike  the  upper  part  of  the  test  tube,  as  the  cork 
may  be  ignited.    At  first  the  heat  expands  the  air  and  a  few 
bubbles  escape  ;  at  a  higher  temperature  the  oxygen  is  evolved. 
Regulate  the  heat  so  as  to  secure  a  uniform  and  not  too  rapid 
evolution  of  the  gas.    By  displacement  of  water  collect  three 

[12] 


or  four  wide-mouthed  bottles  (250-cc.)  of  the  gas.    Before  the 
heat  is  withdrawn  remove  the  cork  from  the  tube  (why?). 


FIG.  16 

Reserve  the  tube  with  its  contents  for  study  in  §  19.  What 
is  the  source  of  the  oxygen  ?  What  is  the  function  of  the 
manganese  dioxide  ? 
16.  Preparation  from 
sodium  peroxide  and 
water.  Sodium  per- 
oxide is  a  white  solid 
containing  41  per 
cent  of  oxygen,  and 
when  treated  with 
water,  a  part  of  this 
is  set  free.  Arrange 
an  apparatus  accord- 
ing to  Fig.  17.  By 
means  of  a  short 
piece  of  rubber  tub- 
ing A  connect  the 
funnel  B  with  a  glass 

tube  (7,  pinching  the  tube  shut  with  a  screw  clamp.    Place 
about  5  g.  of  sodium  peroxide  in  the  bottom  of  D  and  partly 

[13] 


FIG.  17 


^^^ 

^HdU-^ 

""""^ — 7? 
7&M    xfe^fc 


^inju^^L^  ^Jt*^ '  s3Z  suu^&f 


^7 


fill  the  funnel  with  warm  water.  Very  cautiously  open  the 
screw  clamp  so  that  the  water  will  run  down  and  fall,  drop 
by  drop,  upon  the  peroxide.  A  steady  current  of  oxygen  is 
given  off,  and  the  gas  should  be  collected  as  in  the  preced- 
ing exercise.  (A  separatory  funnel  may  be  used  to  advantage 
instead  of  the  funnel  and  screw  clamp.) 

17.  Preparation  from  potassium  permanganate  and  hydrogen 
peroxide.    Using  the  apparatus  shown  in  Fig.  17,  place  a  solu- 
tion of  5  g.  of  potassium  permanganate  in  25  cc.  of  water  in 
D  and  add  3  cc.  of  sulfuric  acid.    Fill  the  funnel  with  com- 
mercial hydrogen  peroxide  and  allow  it  to  flow  very  slowly 
into  the  solution  in  D. 

18.  Properties  and  conduct  of  oxygen,    a.  Note  the  prop- 
erties of  the  gas.    (The  slight  cloud  often  present  when  oxy- 
gen is  prepared  from  potassium  chlorate  is  due  to  an  impurity 
and  will  disappear  if  the  gas  is  allowed  to  stand  over  water.) 

b.  Repeatedly  thrust  a  glowing  splint  into  a  bottle  of  the  gas. 

c.  Heat  some  sulfur  in  a  deflagrating-spoon  until  ignited. 
Note  the  color  and  size  of  the  sulfur  flame.    Now  lower  the 
burning  sulfur  into  a  bottle  of  oxygen  and  note  the  change. 

d.  Tip  a  piece  of  picture-frame  wire,  12  or  15  cm.  long, 
with  sulfur  by  wrapping  a  bit  of  cotton  about  the  end  of 
the   wire    and  dipping  this    into    melted  sulfur.    Ignite  the 
sulfur  by  holding  it  in  a  Bunsen  flame  for  an  instant,  then 
thrust  the  wire  into  a  bottle  of  oxygen. 

Are  the  changes  observed  in  &,  c,  and  d  examples  of  chem- 
ical action?  What  becomes  of  the  oxygen?  What  is  the 
name  of  the  product  formed  in  each  case  ?  | 

19.  Separation  of  the  compounds  present  in  the  residue  left 
in  the  preparation  of  oxygen.    Heat  the  tube  containing  the 
residue  obtained  in  §  15  until  no  more  oxygen  is  evolved. 
After  the  tube  is  cool,  half  fill  it  with  water  and  shake  the 
contents  thoroughly.    After  a  few  minutes  filter  off  the  solid 
matter  (what  is  it?).     Evaporate  the  filtrate  to  a  volume 
of  4  or  5  cc.,  and  set  it  aside  until  crystals  are  deposited. 
Convince  yourself  that  the   substance  is  different  from  the 
potassium  chlorate  with  which  you  started. 

[14] 


CHAPTER  III 
HYDROGEN 

20.  Preparation  of  hydrogen  from  water.    Fill  a  test  tube 
with  water  and  invert  it  in  a  beaker  of  water.    Wrap  a  piece 
of  sodium,  the  size  of  a  small  pea,  in  a  bit  of  filter  paper  pre- 
viously moistened  with  coal  oil.    Raise  the  inverted  test  tube 
until  its  mouth  dips  just  below  the  surface  of  the  water  in 
the  beaker,  and  quickly  insert  the  sodium.    Stand  at  arm's 
length,  as  a  slight  explosion  sometimes  occurs.    Notice  that 
the  sodium  decomposes  the  water,  liberating  a  gas  which  is 
caught  in  the  tube.    Holding  the  tube  mouth  downward,  test 
the  gas  by  quickly  inserting  a  burning  splint.    Does  the  gas  act 
(ike  oxygen  ?    What  is  the  source  of  it  ?    What  other  meth- 
ods may  be  employed  for  obtaining  it  from  the  same  source  ? 

21.  Preparation  from  sodium  hydroxide.    Provide  a  hard- 
glass  test  tube  with  a  stopper  and  a  delivery  tube,  as  shown 
in  Fig.  16.    Finely  powder  about  2g.  of  solid  sodium  hydrox- 
ide in  a  mortar  and  mix  it  with  a  little  more  than  an  equal 
weight  of  zinc  dust.    Place  the  mixture  in  the  bottom  of  the 
test  tube,  clamping  the  tube  in  a  nearly  horizontal  position, 
the  bottom  being   slightly  higher  than  the  mouth  (why  ?). 
Heat  the  mixture  carefully,  collecting  the  gas  evolved  in  a 
test  tube  by  displacement  of  water.    Test  it  with  a  lighted 
splint  as  in  §  20.    What  is  the  source  of  the  gas  ? 

22.  Preparation  from  acids,     a.  In  clean  test  tubes  place 
samples  of  the  following  metals :  iron,  tin,  lead,  magnesium, 
aluminium,  zinc,  and  copper.    Prepare  some  dilute  hydrochlo- 
ric acid  by  adding  about  20  cc.  of  the  concentrated  reagent 
to  an  equal  volume  of  water.    Pour  4  or  5  cc.  of  the  dilute 
acid  in  each  of  the  test  tubes,  gently  warming  the  solution 

[15] 


I 


if  no  action  is  noticed.  In  which  cases  do  you  observe  the 
formation  of  gas  bubbles  ?  Where  do  they  form  ?  Do  not 
mistake  bubbles  of  air  or  steam  for  the  gas.  How  can  you 
tell  the  difference  ?  Does  the  condition  of  the  metal  make 
any  difference  (compare  the  action  of  the  acid  on  a  tack 
with  its  action  on  iron  filings)  ?  Do  all  of  the  metals  yield 
hydrogen  with  acids  ? 

b.  Repeat  the  experiment,  replacing  the  hydrochlorrc  acid 
with  nitric  acid  diluted  with  an  equal  volume  of  water.  In 
what  cases  is  a  gas  evolved  ?  What  differences  do  you  note 
as  compared  with  a  ?  In  each  case  test  the  gas  with  a  lighted 


FIG.  18 

splint.  Is  it  hydrogen  ?  Do  the  metals  which  give  hydrogen 
with  hydrochloric  acid  give  it  with  all  acids  ? 

c.  Select  two  test  tubes  and  in  each  place  a  few  pieces  of 
mossy  zinc,  adding  enough  water  to  cover  the  metal.  To  one 
test  tube  add  a  few  drops  of  a  solution  of  copper  sulfate. 
What  change  do  you  notice  ?  Now  add  to  each  about  1  cc. 
of  hydrochloric  acid.  Which  evolves  hydrogen  the  more 
rapidly  ?  What  is  the  function  of  the  copper  sulfate  (§  14)  ? 

23.  Usual  laboratory  method.  Arrange  a  hydrogen  genera- 
tor according  to  Fig.  18,  in  which  D  represents  a  wide-mouthed 
bottle  of  about  250-cc.  capacity.  The  gas  delivery  tube  B,  C  is 
the  same  as  that  used  in  the  preparation  of  oxygen  (Fig.  16). 
The  funnel  tube  A  must  extend  nearly  to  the  bottom  of 
the  bottle  (why?).  Put  10  g.  of  mossy  zinc  (why  mossy 

[16] 


1 


/^n.0  3  ^^ 


..X^-fi-'fa-t 


zinc  ?)  into  D  and  add  a  few  drops  of  a  solution  of  cop- 
per sulfate.  Pour  just  enough  water  through  the  funnel 
tube  to  cover  the  zinc.  Prove  that  the  apparatus  is  air-tight 
by  blowing  into  the  delivery  tube  until  the  water  is  forced 
nearly  to  the  top  of  the  funnel  tube ;  then  quickly  close  the 
rubber  tube  either  by  tightly  pinching  it  or  by  placing  the 
tongue  firmly  against  its  end.  If  the  apparatus  is  air-tight, 
the  water  in  the  funnel  tube  will  not  fall  (why  ?).  Prepare 
some  dilute  sulfuric  acid  by  sloivly  pouring  15  cc.  of  concen- 
trated acid  into  a  beaker  containing  50  cc.  of  water.  Stir  the 
water  with  a  glass  rod  while  the  acid  is  being  added.  Notice 
that  the  acid  is  poured  into  the  water  —  never  the  reverse.  Cool 
the  mixture  and  pour  a  few  drops  of  it  through  the  funnel  tube. 
Hydrogen  is  at  once  evolved.  Enough  of  the  acid  must  be 
added  from  time  to  time  to  cause  a  gentle  and  continuous 
evolution  of  the  gas.  It  is  evident  that  the  first  gas  which 
passes  over  is  a  mixture  of  hydrogen  and  air.  The  student 
must  remember  that  such  a  confined  mixture  of  hydrogen  and  air 
or  hydrogen  and  oxygen,  if  ignited,  explodes  with  great  violence. 
On  this  account  see  that  the  end  of  the  delivery  tube  is  not 
brought  near  any  flame.  Determine  when  the  hydrogen  is 
free  from  air  by  collecting  test  tubes  full  of  gas  and  igniting 
it,  holding  the  tube  mouth  downward.  If  pure,  the  gas  burns 
quietly,  otherwise  there  is  a  slight  explosion.  After  all  the 
air  has  been  expelled  from  the  generator,  collect  four  bottles 
(250-cc.,  wide-mouthed)  of  the  gas.  What  is  the  source  of 
the  hydrogen  ?  What  is  the  use  of  the  zinc  ?  Why  is  the 
copper  sulfate  solution  added  ?  Remove  the  stopper  from 
the  generator,  add  a  few  more  fragments  of  zinc,  and  set 
the  generator  aside  until  the  gas  ceases  to  be  evolved.  Suffi- 
cient zinc  should  be  used  so  that  at  least  a  small  portion  of 
it  remains  undissolved.  Filter  the  liquid  from  the  undissolved 
zinc  into  an  evaporating-dish  and  evaporate  it  to  dryness. 
What  is  left  in  the  dish  ?  After  the  dish  is  cool,  dissolve  the 
residue  in  as  little  hot  water  as  possible  and  set  it  aside 
until  crystals  are  deposited.  Compare  these  with  crystals 
of  zinc  sulfate. 

[17] 


x— X — 


,: 
5  o  y, 


24.  Properties  of  hydrogen,  a.  Thrust  a  lighted  splint  into 
a  bottle  of  the  gas  held  mouth  downward.  Slowly  withdraw 
the  splint  and  again  thrust  it  into  the  gas.  Describe  the 
results.  What  do  they  prove  ? 

b.  Fill  a  small  (60-cc.)  wide-mouthed  bottle  or  a  test  tube 
one  third  full  of  water,  and  invert  it  in  a  pneumatic  trough. 
Displace  the  remaining  water  with  hydrogen.    What  does  the 
bottle  now  contain  ?  Withdraw  it  from  the  water,  and,  hold- 
ing it  at  arm's  length,  quickly  bring  it  mouth  downward  over 
a  flame.    What  do  the  results  proA^e  ? 

c.  Uncover  a  bottle  (mouth  upward)  of  the  gas.     After 
one  minute  test  for  the  presence  of  hydrogen  with  a  lighted 
splint.   Repeat,  keeping  the  bottle  mouth  downward.   Describe 
the  results.    Is  the  gas  heavier  or  lighter  than  air  ? 

d.  Wash  out  the  hydrogen  generator,  then  remove  the  rub- 
ber delivery  tube  and  connect  the  tube  A  (Fig.  19)  with  the 
drying-tube  B  by  a  short 


O 


piece  of  rubber  tubing. 
The  tube  B  is  filled  with 
pieces  of  calcium  chloride 
(why  ?),  held  in  place  by 
loose  plugs  of  cotton  at 
each  end  of  the  tube. 
The  outer  end  of  the  bent 
glass  tube  C  is  drawn  to 
a  jet.  After  the  appara- 
tus has  been  approved 
by  the  instructor,  charge 
the  generator  with  6  or  8  g. 
of  zinc,  add  a  few  drops  of  a  solution  of  copper  sulfate,  cover 
the  zinc  with  water,  and  add  dilute  sulfuric  acid  as  in  §  23. 
After  all  the  air  has  been  expelled  (prove  this  by  slipping  a 
piece  of  rubber  tubing  over  the  tube  C  and  testing  samples 
of  the  gas  collected  over  water)  wrap  a  towel  carefully  about 
the  generator  and  cautiously  ignite  the  hydrogen.  Test  the 
heat  of  the  flame  by  holding  in  it  different  objects,  such  as  a 
splint,  a  piece  of  picture-frame  wire,  a  bit  of  charcoal. 

[18] 


FIG.  19 


w- 


K1 


•     c        f    •     • 

-h  "^r3?L-i 

•  .••/1       /^t>^W"*"V 


<A- 

4^.,-a, 


l^c/fc-  > 


jj^7cJ^*ky^t- 


25.  The  oxyhydrogen  blowpipe.    If  the  instrument  is  avail 
able,  examine  the  structure   of  the  oxyhydrogen   blowpipe. 
Draw  a  diagram  representing  a  cross  section  of  it.    Compare 
it  with  the  ordinary  laboratory  blowpipe  (blast  lamp).    Why 
not  have  a  short  inner  tube  instead  of  a  long  one  ? 

26.  Qualitative  synthesis  of  water.    Generate  hydrogen  as 
in  §  24,  d  (note  precautions).    Ignite  the  hydrogen  and  hold 
over  the  flame  a  cold,  dry  beaker.    What  substance  condenses 
on  the  sides  of  the  beaker  ?    Account  for  its  formation.    The 
method  of  formation  proves  the  presence  of  what  elements 
in  this  substance  ? 

27.  Reduction.    Arrange  an  apparatus  according  to  Fig.  20, 
in  which  a  hard-glass  tube  (7,  about  30  cm.  long  and  from  8 


o 


FIG.  20 

to  10  mm.  in  diameter,  is  substituted  for  the  glass  jet  C  of 
Fig.  19.  Spread  1  g.  of  black  copper  oxide  in  a  layer  in  the 
tube  10  or  12  cm.  from  the  end  which  is  connected  with  the 
drying-tube.  This  is  done  as  shown  in  Fig.  15,  except  that 
the  tube  must  be  held  in  a  horizontal  position  and  the  oxide 
deposited  at  the  proper  place  by  turning  the  paper  over. 
Now  generate  hydrogen  as  in  §  23.  After  all  the  air  has 
been  expelled  from  the  apparatus  and  the  generator  has  been 
wrapped  in  a  towel,  cautiously  heat  the  copper  oxide  to 

[19] 


uij^^L  ^4f^ 

J^sf*    ll.^t^T^ 


redness  with  a  "  wing-top  "  burner.  Note  the  condensation  of 
moisture  in  the  cold  portions  of  the  tube.  Account  for  its 
formation.  What  change  has  the  copper  oxide  undergone  ? 
Is  there  any  visible  evidence  of  this  change  ?  Suggest  a 
method  of  finding  the  weight  of  the  water  formed. 

28.  Oxidation,    a.  Without  removing  the  contents  of  the 
hard-glass  tube,  replace  the  hydrogen  generator  and  drying, 
tube  by  the  oxygen  generator. described  in  §  15.    Pass  a  slow 
current  of  oxygen  through  the  tube,  gradually  heating  its 
contents  to  redness.    What  change  takes  place  ?    How  does 
the  product  compare  with  the  original  copper  oxide  ? 

b.  Explain  the  terms  reduction,  oxidation,  reducing  agent, 
and  oxidizing  agent,  and  give  examples  of  each  from  §  27 
and  §  28,  a. 

29.  Measurement  of  gas  volumes,    a.  Partly  fill  a  graduated 
tube  A  (Fig.  21)  with  water  and  invert  it  in  a  vessel  of  water 
so  that  the  level  of  the  liquid  in  the  tube 

is  above  that  of  the  liquid  in  the  vessel. 
Clamp  it  in  this  position  and  read  the  vol- 
ume of  the  inclosed  air.  What  conditions 
affect  the  volume  of  a  gas  ?  What  are 
the  conditions  adopted  as  standard  in  the 
measurement  of  gases  ?  From  the  volume 
as  measured  above,  calculate  the  volume 
under  these  conditions. 

&.  Adjust  the  tube  so  that  the  level 
within  is  some  distance  below  the  level 
without.  Read  the  new  volume  and  re- 
duce to  standard  conditions.  If  the  liquid 
employed  had  been  mercury  instead  of 
water,  what  changes  would  this  make  in  the 
calculation  ?  The  correction  for  aqueous  tension  is  explained 
n  pages  91,  92  of  the  text  and  should  be  studied  at  this  point. 

30.  The  law  of  Gay-Lussac  (or  of  Charles).    The  expansion 
of  a  definite  volume  of  air  through  a  definite  interval  of  tem- 
perature may  be  measured  as  follows :  Provide  an  Erlenmeyer 
or  a  Florence  flask  of  about  200-cc.  capacity,  A  (Fig.  22),  and 

[20] 


00 


FIG. 21 


1 


fit  it  with  a  rubber  stopper  carrying  a  glass  tube  B,  about 
8  cm.  in  length  and  not  too  small  in  bore.  When  tightly  in- 
serted, the  stopper  should  be  about  halfway  in  the  flask,  as 
shown  in  Fig.  22.  Immerse  the  dry,  empty  flask  in  a  beaker 
of  water,  weighting  it  down  with  a  ring  clamp  C,  as  shown  in 
the  figure,  or  holding  it  in  place  by  clamping  the  glass  tube. 
Heat  the  bath  at  the  boiling  point  of  water  for  at  least  fifteen 
minutes,  when  it  may  be  assumed  that  the  air  in  the  flask  is 
also  at  that  temperature  (^2). 

Press  a  finger  tip  firmly  against  the  end  of  the  tube,  quickly 
lift  the  flask  from  the  bath,  using  a  towel  or  a  test-tube  clamp 
to  protect  the  fingers,  and  dip  the  end  of  the  tube  below  the 
surface  of  water  in  a  pneumatic 
trough.  Now  remove  the  finger  from 
the  end  of  the  tube  and  allow  water 
from  the  tap  to  run  on  the  bottom 
of  the  flask  until  the  air  in  the  flask 
has  cooled  to  the  temperature  of  the 
tap  water.  Note  this  temperature 
(^).  The  volume  of  water  enter- 
ing the  flask  is  equal  to  the  con- 
traction in  the  original  volume  (  F2) 
of  air  in  cooling  from  T2  to  T^ 

Adjust  the  level  within  and  without  the  flask,  close  the 
end  of  the  tube,  and  remove  the  flask,  placing  it  right  side 
up  on  the  table.  Wipe  it  dry,  paste  a  paper  label  to  mark 
the  level  of  the  lower  end  of  the  stopper  (or  mark  it  with  a 
rubber  band),  and  cautiously  remove  the  stopper,  letting  any 
water  remaining  in  the  tube  run  into  the  flask.  Carefully 
pour  the  water  into  a  graduated  cylinder  and  note  its  volume 
as  accurately  as  possible,  calling  this  a.  Fill  the  flask  with 
water,  pushing  in  the  stopper  to  the  mark  and  taking  care 
that  the  tube  is  also  full  and  that  there  is  no  air  bubble 
under  the  stopper.  Carefully  remove  the  stopper,  allowing 
the  tube  to  drain  into  the  space  vacated  by  the  stopper. 
Lastly,  measure  the  water  in  a  graduated  tube,  calling  this 


FIG.  22 


volume  Fa. 


[21] 


Now  F"2  —  a  is  the  volume  of  air  at  Tl  (volume  Ft)  which 
will  just  fill  the  flask  and  tube  (volume  F2)  at  1\  (100°  C., 
or  373°  A.).  In  other  words,  in  the  interval  T2  —  Tl  the  gas 
contracts  from  F2  to  Fr  Calculate  the  volume  that  F2  ought 
to  give  at  T^  assuming  the  accuracy  of  the  law  of  Gay-Lussac. 
Compare  the  value  found  by  calculation  with  the  experi- 
mental value.  What  is  the  percentage  of  error  ?  What  sources 
of  error  can  you  think  of  ? 

31.  Percentage  of  oxygen  in  potassium  chlorate;  weight  of 
i  liter  of  oxygen.  Prepare  the  apparatus  shown  in  Fig.  23. 
A  represents  the  hard-glass  test  tube  used  in  the  preparation 
of  oxygen,  B  is  a  common  narrow-mouthed  bottle,  having  a 
capacity  of  about  1  liter.  The 
rubber  tube  C  is  provided  with 
a  screw  clamp  D  for  closing  the 
tube,  and  ends  in  a  piece  of  glass 
tubing  drawn  out  to  a  jet,  the 
internal  diameter  of  the  jet  being 
about  2  mm.  The  bottle  is  nearly 
filled  with  water,  as  shown  in  the 
figure,  and  allowed  to  stand  until 
it  acquires  the  room  temperature. 
The  tube  A  is  now  removed,  and 
a  gentle  suction  is  applied  to  the  glass  jet,  at  the  end  of 
tube  C.  The  water  siphons  over  through  the  tube  C  into  the 
beaker  and  is  allowed  to  run  for  a  moment  to  fill  completely 
both  the  rubber  tube  and  the  glass  tube.  The  rubber  tube  is 
then  quickly  closed  with  the  screw  clamp. 

Now  thoroughly  clean  and  dry  the  tube  A,  and  carefully 
weigh  it ;  then  introduce  about  1  g.  of  potassium  chlorate  into 
the  bottom  of  the  tube  by  means  of  a  folded  paper  (Fig.  15), 
and  reweigh.  Attach  the  tube,  as  shown  in  Fig.  23,  care 
being  taken  to  have  the  apparatus  air-tight.  The  pressure  of 
the  air  within  the  bottle  is  now  adjusted  to  that  of  the  air 
outside,  as  follows:  Water  is  added  to  the  beaker  (if  neces- 
sary) until  the  end  of  the  glass  tube  is  covered.  The  screw 
clamp  is  then  opened  and  the  beaker  at  once  raised  vertically 

[22] 


FIG.  23 


until  the  water  in  the  beaker  is  at  the  same  level  as  the  water 
in  the  bottle ;  then  the  screw  clamp  is  closed.  Empty  the 
beaker,  and  return  it  to  the  position  shown  in  Fig.  23.  Now 
open  the  screw  clamp  and  apply  a  gentle  heat  to  the  potassium 
chlorate.  Oxygen  is  evolved  and  forces  the  water  from  the 
bottle  into  the  beaker.  Gradually  increase  the  heat,  and  con- 
tinue the  heating  until  all  the  oxygen  has  been  expelled.  Let 
the  apparatus  stand  until  it  has  acquired  room  temperature, 
care  being  taken  that  the  glass  jet  is  kept  below  the  surface 
of  the  water  in  the  beaker  (why  ?).  Now  bring  the  level 
of  the  water  in  the  beaker  to  that  of  the  water  left  in  the 
bottle,  and,  while  holding  it  in  this  position,  close  the  screw 
clamp  (why  ?).  Carefully  measure  the  water  in  the  beaker ; 
also  take  the  readings  of  the  thermometer  and  the  barometer. 
Disconnect  the  tube  A,  and  carefully  re  weigh  the  tube  and 
its  contents.  Insert  the  values  in  the  following  table : 

Weight  of  tube  A 

Weight  of  tube  A  +  the  potassium  chlorate 

Weight  of  tube  A  +  the  potassium  chloride 

Volume  of  water  in  beaker  =  volume  of  oxygen  evolved     .     . 

Temperature  of  water 

Barometric  reading 


From  your  results  calculate  (1)  the  percentage  of  oxygen 
in  potassium  chlorate,  and  (2)  the  weight  of  1  liter  of  oxygen 
under  standard  conditions.  Compare  your  results  with  those 
given  in  the  text. 

If  sufficient  heat  has  not  been  applied  to  expel  all  the  oxy- 
gen from  the  potassium  chlorate,  incorrect  results,  of  course, 
would  be  obtained  for  the  percentage  of  oxygen  present.  This 
would  have  no  effect,  however,  upon  the  determination  of  the 
weight  of  1  liter  of  the  gas.  It  is  well,  therefore,  after  the 
tube  A  and  the  residue  have  been  accurately  weighed,  to 
reheat  the  tube  and  reweigh.  This  process  should  be  repeated 
until  the  tube  and  contents  suffer  no  loss  of  weight  on  heating. 


[23] 


CHAPTER  IV 


WATER  AND  HYDROGEN  PEROXIDE 

32.  Distillation.    Connect  a  500-cc.  flask  A  with  a  Liebig 
condenser    B    (obtained    from    the    general    storeroom),    as 
shown  in  Fig.  24.    The  tube  C  is  connected  with  the  water 
pipe  by  means  of  rubber  tubing,  and  a  current  of  cold  water 
is  allowed  to  flow  through  the  outer  tube  of  the  condenser. 
Why  is  cold  water  admitted  at  C  rather  than  at  Z>?    Half 
fill  the  flask  with  hy- 

drant water  and  boil 

till  100  cc.  or  more  of 

liquid  has  collected  in 

E,  the  receiver.    Com- 

pare the  distillate  (dis- 

tilled water)  with  the 

hydrant  water  in  ap- 

pearance   and    taste. 

To  what  is  the  differ- 

ence due  ?    Interrupt 

the  process  of  distilla- 

tion and  add  a  few  crystals  of  a  highly  colored  salt  (potassium 

permanganate)  to  the  water  in  the  distilling-flask.    Continue 

the  distillation  of  the  water.   Is  the  distillate  colored  ?    Place 

4  or  5  drops  of   the  distilled  water  on  a  watch  glass  and 

evaporate.    Is  there    any  residue  ?    Repeat,   using   hydrant 

water.    Why  is  distilled  water  used  in  the  laboratory? 

33.  Filtration,  a.  To  about  25  cc.  of  water  in  a  beaker  add 
1  or  2  cc.  of  ammonia  water  and  stir.    Note  the  odor  of  the 
solution.    Add  as  much  bone  black  or  powdered  charcoal  as 
you  can  pile  on  a  five-cent  piece,  and  thoroughly  stir.    Filter 

[24] 


FlG 


the  solution  and  note  the  odor  of  the  filtrate.    How  does  it 
compare  with  that  of  the  original  solution  ? 

b.  To  about  100  cc.  of  water  add  enough  blue  litmus  solu- 
tion to  give  a  decided  blue  color.  Half  fill  a  test  tube  with 
the  solution  and  set  it  aside.  To  the  remainder  add  5  cc.  of  a 
solution  of  alum  and  an  equal  volume  of  dilute  ammonia.  Note 
the  formation  of  a  white  precipitate.  Does  it  remain  entirely 
white  ?  Does  the  solution  remain  as  blue  as  at  first  ?  Stir  the 
solution  for  a  time  and  pour  some  of  it  on  a  filter,  catching 
the  filtrate  in  a  test  tube,  until  the  latter  is  half  full.  Com- 
pare the  color  of  the  filtrate  with  that  of  the  original  solution. 

34.  Solid  residue,    a.  Nearly  fill  ^  y 
a  weighed  evaporating-dish  with  a 

measured  volume  of  water  (not  dis- 
tilled), evaporate  to  dryness,  and 
reweigh.  From  your  results  calcu- 
late the  weight  of  solid  matter  in 
a  liter  of  water. 

b.  Accurately  weigh  a  small 
evaporating-dish  and  a  glass  stir- 
ring-rod. Introduce  about  10  cc. 
of  milk  ancl  reweigh.  Evaporate  to 
dryness  on  a  water  bath  (Fig.  25), 
stirring  the  contents  often  with 
the  glass  rod.  When  the  liquid  has  entirely  evaporated, 
remove  the  beaker  from  the  bath,  cool,  and  weigh.  From 
your  results  determine  the  percentage  of  residue  in  the 
sample  of  milk.  Vinegar  may  be  used  instead  of  milk. 

35.  Chemical  conduct  of  water,    a.  What  is  the  effect  of  a 
moderate  degree  of  heat  upon  water?    Under  what  circum- 
stances may  water  be  decomposed  by  heat? 

b.  Action  upon  oxides.  Dip  a  strip  of  blue  litmus  paper 
into  pure  water.  Is  there  any  change  in  color  ?  Repeat  with 
red  litmus  paper.  Acids  turn  blue  litmus  red,  while  bases 
turn  red  litmus  blue.  Has  water  either  acid  or  basic  proper- 
ties ?  In  a  test  tube  shake  a  small  piece  of  lime  (the  size  of  a 
pea)  with  about  5  cc.  of  water.  Test  the  solution  with  litmus 

[25] 


FIG.  25 


-  Ir 


V^**7  ^c^f^AjJl^    ^r>u>Cc 


O^ 


^7^u^on^O-<-^CC 


"/  * 

^^^  >£<&»- 


(Li*~*.  s. 


^^^£^-z 

fr+^**~ 


*****  /u^tr 

^^tf** 


*rt~&    , 


paper  of  both  colors.  What  inference  do  you  draw?  Set 
fire  to  a  piece  of  sulfur  supported  on  a  deflagrating-spoon  and 
lower  it  into  a  bottle  containing  a  little  water.  After  it  has 
burned  for  a  time,  remove  the  spoon,  cover  the  mouth  of  the 
bottle  with  the  palm  of  the  hand,  and  shake  the  bottle  vigor- 
ously. Test  the  water  with  litmus  paper.  What  conclusions 
do  you  draw  ?  When  water  acts  upon  an  oxide,  what  classes 
of  substances  may  result  ? 

36.  Water  of  crystallization,  or  water  of  hydration.    a.  Heat 
some  small  crystals  of  zinc  sulfate  in  a  dry  test  tube.    What 
evidence  have  you  of  the  presence  of  water  in  the  crystals  ? 
Examine  the  residue.    How  does  it  differ  from  the  original 
crystals  in  appearance  and  in  composition  ? 

b.  Select  some  small  crystals  of  copper  sulfate.    Are  they 
dry?    Fill  a  test  tube  one  fourth  full  of  these  crystals,  and 
heat  as  in  a.    Compare  the  residue  with  the  original  crystals 
in  form,  color,  and  composition.    Dissolve  the  residue  in  as 
little  hot  water  as  possible,  pour  the  solution  (note  the  color 
of  it)  into  an  evaporating-dish,  and  set  aside  until  crystals 
are   deposited.     How   do   these    compare  with   the    original 
crystals   of  copper  sulfate?  — ui 

Does  water  seem  to  be  necessary  to  the  existence  of  zinc 
sulfate  and  copper  sulfate  in  the  usual  crystalline  form  ? 
What  is  such  water  called?  Is  it  combined  with  the  sub- 
stance composing  the  crystal  or  simply  mixed  with  it  ?  Give 
reason  for  your  answer.  Do  all  crystals  contain  it?  (Try 
a  crystal  of  potassium  dichromate.) 

c.  Expose  a  clear  crystal  of  sodium  sulfate  to  the  air  for 
one  or  two  hours.    Notice  the  change  in  its  appearance.    To 
what  is  the  change  due  ?   What  are  such  substances  called  ? 

d.  Place  a  small  piece  of  calcium  chloride  or  of  calcium  ni- 
trate on  a  watch  glass  and  expose  it  to  the  air  for  two  hours 
or  longer.    Note  the  change.   What  term  is  applied  to  bodies 
that  undergo  this  change  ? 

37.  Percentage  of  water   of   hydration   in   copper   sulfate/ 
Accurately  weigh  a  porcelain  crucible  and  cover.   Then  add  2 
or  3  g.  of  small  crystals  of  copper  sulfate  and  again  accurately 

[26] 


;£L^>:  cu*^~~l- 


weigh.  Place  the  covered  crucible  on  the  pipestem  triangle 
and  heat  with  a  gentle  flame  until  the  color  of  the  crystals 
has  entirely  disappeared.  This  will  require  from  twenty  to 
thirty  minutes.  The  tip  of  the  flame  must  not  quite  touch 
the  crucible.  The  product  is  anhydrous  copper  sulfate. 
When  the  crucible  is  cool,  reweigh.  From  your  results  cal- 
culate the  percentage  of  water  of  hydration  in  the  crystals. 
Compare  your  results  with  those  of  other  students  who  have 
used  different  weights  of  crystals. 

38.   Quantitative  synthesis  of  water.    Arrange  an  appara- 
tus as  shown  in  Fig.  26,  in  which  A  represents  the  hydrogen 

0 


B 


E 


I  m 


FIG.  26 

generator,  and  B  and  D  are  tubes  filled  with  dry  calcium 
chloride.  The  hard-glass  tube  C  and  the  porcelain  boat  E  are 
obtained  from  the  storeroom.  The  tube  is  about  35  cm.  in 
length.  Introduce  about  2  g.  of  black  oxide  of  copper  into 
the  boat,  and  weigh  accurately  to  milligrams.  Introduce  the 
boat  into  the  glass  tube  so  that  the  end  of  the  boat  is  about 
8  cm.  from  the  end  of  the  tube  connected  with  D.  Close  the 
ends  of  tube  D  with  short  pieces  of  rubber  tubing,  one  end 
of  each  being  closed  with  a  small  glass  rod ;  then  weigh  the 
tube  accurately.  Remove  the  rubber  tubes  and  glass  rods, 
carefully  preserving  them  for  use  when  the  tube  is  again 
weighed.  Now  connect  the  apparatus  as  shown  in  the  figure, 
taking  care  to  render  it  air-tight.  How  is  this  determined? 
Generate  hydrogen  slowly,  and  when  the  apparatus  is  free 

[27] 


from  air,  heat  the  boat  very  gently,  using  the  wing-top 
burner.  Gradually  increase  the  heat,  all  the  time  maintain- 
ing the  slow  current  of  hydrogen.  When  the  copper  oxide 
is  reduced,  or  nearly  so,  withdraw  the  heat  but  maintain  the 
current  of  hydrogen  until  the  apparatus  is  cool.  If  any  of 
the  water  formed  remains  condensed  in  the  end  of  tube  (7, 
a  very  gentle  heat  is  cautiously  applied  (the  flame  must  not 
strike  the  tube)  until  it  is  driven  into  the  tube  D. 

When  the  apparatus  has  acquired  the  room  temperature, 
disconnect  A  and  attach  D  (Fig.  26)  to  the  short  bent  glass 
tube  in  bottle  B  of  Fig.  23.  The  bottle  is  filled  with  water, 
and  a  portion  of  it  is  slowly  siphoned  over  through  (7,  D 
(Fig.  23).  In  this  way  a  current  of  air  is  drawn  through 
the  apparatus,  displacing  the  hydrogen.  Finally,  disconnect 
the  apparatus,  and  at  once  close  the  ends  of  the  tube  D 
(Fig.  26)  with  the  rubber  tubes  provided  with  glass  rods. 
Weigh  the  boat  and  contents ;  also  the  tube  D.  From  your 
results  calculate  the  composition  of  water. 

In  order  that  the  experiment  may  be  successful,  it  is  ab- 
solutely essential  that  the  current  of  hydrogen  be  maintained 
constantly  until  the  apparatus  is  disconnected ;  otherwise  the 
water  formed  will  not  pass  over  into  the  tube  D. 

39.  Hydrogen  peroxide,  a.  To  about  50  cc.  of  water  in  a 
suitable  flask  add  3  or  4  g.  of  barium  peroxide.  Keep  the 
flask  cool  with  running  water  while  slowly  adding  dilute  sul- 
furic  acid  until  the  solution  changes  blue  litmus  paper  to 
bright  red.  Filter  off  the  precipitate  (barium  sulfate).  The 
clear  solution  contains  hydrogen  peroxide. 

b.  The  following  is  a  test  for  hydrogen  peroxide:  Obtain 
about  1  cc.  of  starch  paste  (side  shelf)  and  dilute  to  5  cc.    Add 
a  very  small  crystal  of  potassium  iodide  (or  1  cc.  of  solution) 
and  then  add  a  few  drops  of  the  solution  of  hydrogen  peroxide. 
An  intensely  blue  color  is  produced. 

c.  Place  about  2  g.  of  finely  powdered  manganese  dioxide 
in  a  test  tube  and  add  3  or  4  cc.  of  the  solution  of  hydrogen 
peroxide.  What  evidence  do  you  observe  that  a  gas  is  evolved  ? 
Test  it  with  a  glowing  splint.    What  is  it  ? 

[28] 


„  o~ 


J^L 


37; 

Jbuvut&n^« 


(0Ju*«-'**4  J 

t*4U+£jL     *V~ 

u^r  ^M^- 


^c^^^  t 

djubS^-*^ 


>-^  ^rtrr^t-^  >  ^ 

•H*Q~~ 


d.  Repeat,  using  an  equal  bulk  of  powdered  charcoal  in 
place  of  the  manganese  dioxide. 

e.  To  1  or  2  cc.  of  a  hot  dilute  solution  of  lead  nitrate  add 
a  solution  of  hydrogen  sulfide.    A  black  precipitate  of  lead 
sulfide  is  produced.    Boil  a  moment  or  two  and  then  allow 
the  precipitate  to  settle.    Pour  off  the  water,  wash  the  pre- 
cipitate once  with  hot  water,  and  then  add  the  solution  of 
hydrogen  peroxide.    What  change  do  you  note  in  the  color 
of  the  precipitate  ?    How  do  you  account  for  the  change  ? 

/.  Pour  3  or  4  cc.  of  a  solution  of  the  coloring  matter 
called  cochineal  in  a  test  tube.  Add  1  or  2  cc.  of  the  solu- 
tion of  hydrogen  peroxide  and  shake  the  contents  of  the 
tube  vigorously.  Is  the  color  of  the  cochineal  changed? 
What  property  of  hydrogen  peroxide  is  illustrated  by  this 
experiment  ? 

g.  Pour  3  or  4  cc.  of  the  solution  of  hydrogen  peroxide 
into  a  test  tube  and  add  rather  more  than  an  equal  volume 
of  ether.  Shake  the  solution  vigorously  and  notice  that  the 
ether  quickly  rises  to  a  colorless  layer  on  top  of  the  solution. 
Now  add  1  drop  of  a  solution  of  potassium  dichromate  and 
again  shake.  Is  the  ether  layer  still  colorless  ?  A  blue  color 
in  the  ether  is  a  delicate  test  for  hydrogen  peroxide. 


(.29] 


CHAPTER  V 
THE  STATES  OF  MATTER 

40.  Evaporation.  Place  about  1  cc.  of  ether  or  of  chloroform 
on  a  watch  crystal  and  blow  upon  the  surface  of  the  liquid. 
Why  does  it  evaporate  so  fast  ?    Do  you  notice  any  change  in 
the  temperature  of  the  watch  glass  ?    Why  is  this  ? 

41.  Boiling  point.    Pour  about  250  cc.  of  distilled  water  into 
a  suitable  flask  or  beaker,  and  heat  it  at  such  a  rate  that  the 
temperature  rises  slowly  but  steadily.    Use  a  thermometer  to 
keep  the  water  stirred.    At  what  temperature  do  you  observe 
bubbles?    Where  do  they  appear  to  form?    What  are  they? 
At  about  what  temperature  do  larger  bubbles  begin  to  form 
at  the  bottom  of  the  vessel  ?  What  becomes  of  them  ?  Why  ? 
At  what  temperature  do  they  move  freely  up  through  the 
liquid  to  the  surface  ?    Do  they  get  larger  or  smaller  as  they 
rise  ?    Why  ?    When  the  water  is  gently  boiling,  try  increas- 
ing the  heat.    Is  the  boiling  any  more  energetic  ?    Does  the 
temperature  rise  any  ?   How  do  you  define  the  boiling  point  ? 

42.  Freezing  point ;  melting  point,   a.  Powder  about  3  g.  of 
sulfur  and  place  it  in  a  small  test  tube.  Heat  it  very  slowly  a 
little  above  a  small  Bunsen  flame,  stirring  it  with  a  thermom- 
eter.   At  what  temperature  does  the  sulfur  begin  to  melt  ? 
During  the  process  of  melting  watch  the  thermometer"  closely. 
Does  the  temperature  change  during  the  process  ?    When  all 
the  sulfur  has  melted,  remove  the  flame  and  let  the  sulfur 
cool.    At  what  temperature  does  it  begin  to  solidify  ?    Does 
the  temperature  change  during  solidification  ? 

b.  Repeat  the  whole  experiment  with  a  little  paraffin.  Do 
you  notice  any  differences  ?  To  the  melted  paraffin  add  a  little 
vaseline.  How  does  this  affect  the  melting  point  of  the  paraffin  ? 

[30] 


w 


/ 


.  73 


AL^f 


c.  If  the  salts  are  available,  determine  the  melting  point 
of  crystallized  calcium  nitrate,  crystallized  sodium  acetate, 
crystallized  sodium  thiosulfate. 

43.  Sublimation.  Place  2  or  3  g.  of  ammonium  chloride  in 
a  small  evaporating-dish  and  warm 'it  gently.  Does  it  melt? 
What  change  do  you  notice  ?  Cover  the  dish  with  a  small 
inverted  funnel  and  continue  to  heat  very  gently  so  that 
not  too  much  smoke  issues  from  the  stem  of  the  funnel. 
After  a  time  let  the  dish  cool.  Is  there  any  solid  upon 
the  sides  of  the  funnel  ?  How  did  it  get  there  ?  What  is 
such  a  process  called? 


CHAPTER.  VI 

THE  LAWS  OF  CHEMICAL  COMBINATION 
I.   THE  LAW  OF  DEFINITE  COMPOSITION 

44.  Deter mmation  of  the  weight  of  oxygen  which  combines 
with  a  definite  weight  of  magnesium,  a.  Thoroughly  clean 
and  dry  the  porcelain  crucible  and  lid  included  in  your  list 
of  apparatus  and  weigh  them  accurately.  Obtain  a  piece  of 
magnesium  ribbon  weighing  not  to  exceed  0.5  g.  and  scrape 
it  with  a  knife  until  the  surface  is  bright.  Cut  the  ribbon 
into  pieces  1  or  2  cm.  in  length,  place  the  pieces  in  the  cru- 
cible, and  again  weigh.  Now  place  the  covered  crucible  on 
the  triangle  and  apply  a  gradually  increasing  heat.  Magne- 
sium burns  violently  when  freely  exposed  to  the  air,  hence 
the  cover  should  be  left  on  the  crucible  until  the  oxidation  is 
nearly  complete.  This  will  require  about  twenty  minutes. 
Cautiously  jemove  the  cover  by  means  of  the  forceps,  but  hold 
it  just  above  the  crucible  so  that  it  can  be  returned  at  once 
in  4&e  the  magnesium  should  begin  to  burn.  Continue  the 
heating  until  the  substance  no  longer  glows  when  the  cover 
is  removed.  Finally,  tip  the  crucible  partly  on  its  side,  so 
as  to  give  free  access  of  air,  and  apply  a  strong  heat  for  a 
few  minutes  longer.  Withdraw  the  flame  and  re  weigh  the 
crucible,  first  allowing  it  to  cool. 

b.  From  your  results  calculate  the  weight  of  oxygen  which 
combines  with  1  g.  of  magnesium.  Repeat  the  experiment, 
using  a  different  weight  of  'magnesium.  In  case  the  time  is 
not  sufficient  to  repeat  the  experiment,  compare  your  result 
with  those  obtained  by  other  members  of  the  class.  Are  they 
in  accord  with  the  law  of  definite  composition  ?  If  not,  how 
do  you  account  for  their  variation  from  the  law? 

[32] 


^^-  s*****  ^a^e^^A^. 


~     3. 


•  5*5* 


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fin 


-  / 


•  <f 


45.  Determination  of  the  percentage  of  oxygen  in  the  black 
oxide  of  copper.    Use  the  apparatus  shown  in  Fig.  26,  except 
that  a  short  glass  tube  is  substituted  for  the  tube  D.    Accu- 
rately weigh  two  porcelain  boats,  introduce  different  weights 
of  black  copper  oxide,  and  re  weigh.    Place  the  boats  in  the 
glass  tube  and  proceed  with  the  experiment  in   accordance 
with  §  38.  Apply  the  heat  first  to  the  boat  nearest  the  hydro- 
gen generator.    Continue  the  heating  for  fifteen  or  twenty 
minutes,  in  order  completely  to  reduce  the  oxide.    After  the 
apparatus  has  cooled,  weigh  the  boats  and  calculate  the  per- 
centage of  oxygen  in  each  sample  of  the  oxide.    What  do 
the  results  show  ? 

46.  Determination  of  the  weight  of  common  salt  obtained  by 
adding  hydrochloric  acid  to  a  definite  weight  of  sodium  bicarbon- 
ate.   When  hydrochloric  acid  is  added  to  sodium  bicarbonate 
there  are  formed  common  salt  (sodium  chloride),  water,  and 
carbon  dioxide.   To  determine  the  weight  of  salt  formed  from 
a  definite  weight    of    sodium   bicarbonate,   proceed    as    fol- 
lows: Carefully  weigh  the  evaporating-dish  and  watch  glass. 
Transfer  to  the  dish  about  1  g.  of  sodium  bicarbonate  and 
reweigh.    Pour  4  or  5  cc.  of  water  on  the  bicarbonate,  and 
place  the  watch  glass  on  the  dish  so  that  only  the  lip  of  the 
dish  remains,,  uncovered.    Now  pour  down  the  lip  of  the  dish 
2  or  3  drops  of  hydrochloric  acid.    Wait  until  the  efferves- 
cence caused  by  the  escape  of  the  carbon  dioxide  ceases,  then 
add  a  few  drops  more  of  the  acid.    Repeat  until  the  addition 
of  the  acid  no  longer  causes  any  effervescence.    Now  lift  the 
watch  glass  and  with  a  little  water  carefully  rinse  back  into 
the  dish  the  liquid  which  has  collected  on  the  undersurface 
of  the  watch  glass.    Lay  the  watch  glass  aside  and  slowly 
evaporate  the  solution,  taking  care  that  it  does  not  quite  boil 
(why  ?).    When  only  a  few  drops  of  the  liquid  remain,  coyer 
the  dish  with  the  watch  glass  and  increase  the  heat.    The  tip 
of  the  flame  should  just  touch  the  bottom  of  the  dish.    Con- 
tinue the  heating  until  there  is  no  more  liquid  left  in  the  dish 
or  clinging  to  the  undersurface  of  the  glass.    Then  withdraw 
the  heat,  and  after  the  dish  is  cool,  reweigh.    The  residue  is 

[33] 


ID 


" 


-j^.^^C^f^^- 


y 5-7, 1  x  -Y0  —  7771    ^'? 


common  salt.  From  your  results  calculate  the  weight  of  salt 
formed  from  1  g.  of  the  bicarbonate.  Repeat  the  experiment, 
using  a  different  weight  of  sodium  bicarbonate.  Compare  the 
results  of  the  experiments.  In  case  the  time  is  not  sufficient 
to  repeat  the  experiment,  compare  your  result  with  those 
obtained  by  other  members  of  the  class.  Are  they  in  accord 
with  the  law  of  definite  composition  ? 

II.   THE  LAW  OF  MULTIPLE  PROPORTION 

47.  Determination  of  the  percentage  of  oxygen  in  each  of  the 
two  oxides  of  copper.    Repeat  §  45,  substituting  the  red  oxide 
of  copper  for  the  black  oxide  in  one  of  the  boats.    From  your 
results  calculate  the  weight  of  oxygen  combined  with  1  g.  of 
copper  in  each  of  the  compounds.    Are  the  results  in  accord 
with  the  law  of  multiple  proportions  ? 

48.  The  determination  of  the  weight  of  oxygen  in  potas- 
sium chlorate  and  in  potassium  perchlorate.    Determine  the 
weight  of   oxygen    in   potassium   perchlorate    by  heating  a 
weighed  portion  in  the  hard-glass  test  tube  used  in  the  prepa- 
ration of  oxygen.   Compare  your  results  with  those  obtained  in 
heating  potassium  chlorate  (§  31).    Both  potassium  chlorate 
and  potassium  perchlorate  are  compounds  of  potassium,  chlo- 
rine, and  oxygen.    When  either  of  these  compounds  is  heated 
sufficiently,  the  oxygen  is  entirely  expelled,  and  potassium 
chloride  is  left.     From  your  results  calculate  the  weight  of 
oxygen  combined  with  1  g.  of  potassium  chloride  in  each  of 
the  two  compounds. 

III.    THE  LAW  OF  COMBINING  WEIGHTS 

49.  Determination  of  the  volume  of  hydrogen  displaced  by  a 
definite  weight  of  zinc.    Arrange  an  apparatus  in  accordance 
with  Fig.  27.    Here  A  is  a  60-cc.  wide-mouthed  bottle.    The 
bottle  B  and  the  accompanying  tubes  Z>,  E  are  the  same  as 
shown  in  Fig.  23,  except  that  the  hard-glass  test  tube  has  been 
removed.   The  glass  tube  F  is  drawn  out  to  a  jet  at  the  lower 
end,  and  the  other  end  is  connected  to  the  funnel  by  means  of 

.    [34] 


a  rubber  tube,  which  can  be  closed  by  the  screw  clamp  C.  Dis- 
connect the  bottles  A  and  B  at  D  and  fill  the  bottle  B  and  the 
exit  tube  with  water,  as  in  §  31.  Weigh  out  accurately  about 
1  g.  of  zinc ;  then  remove  the  cork  from  bottle  A  and  trans- 
fer the  zinc  to  the  bottle.  Also  add  to  the  bottle  1  drop  of 
a  solution  of  copper  sulfate.  Close  the  screw  clamp  (7,  and 
nearly  fill  the  funnel  with  a  dilute  solution  of  sulfuric  acid 
prepared  as  directed  in  §  23. 
Place  a  beaker  under  the 
tube  F  and  open  the  screw 
clamp  until  both  the  rubber 
and  glass  tubes  are  com- 
pletely filled  with  the  dilute 
acid  ;  then  quickly  close  the 
clamp.  Connect  the  appa- 
ratus just  as  shown  in  the 
figure,  care  being  taken  to 
make  the  joints  air-tight. 
Now  adjust  the  pressure  of 
the  air  inside  the  bottles, 
as  directed  in  §  31.  After 

opening  the  clamp  E,  partially  open  the  clamp  C,  and  allow  a 
few  drops  of  the  acid  to  flow  into  the  bottle  A.  As  the  hydro- 
gen is  evolved  it  forces  the  water  from  bottle  B  into  the  beaker. 
More  acid  is  added  from  time  to  time.  After  the  zinc  has  all 
dissolved  and  the  apparatus  has  acquired  the  room  tempera- 
ture, again  adjust  the  pressure  of  the  gas  within  the  bottles  and 
close  the  clamp  E.  Insert  the  values  in  the  following  table : 


FIG. 27 


Weight  of  zinc  taken  . 

Volume  of  water  forced  into  the  beaker 
Volume  of  liquid  left  in  bottle  A  . 
Volume  of  hydrogen  liberated  . 

Temperature 

Barometric  reading 


Reduce  to  standard  conditions  the  volume  of  the  hydrogen 
obtained  and  compare  it  with  the  theoretical  results.  What 
sources  of  error  are  involved  in  the  experiment  ? 

[35] 


CHAPTER  VII 


CARBON  AND  CARBON  DIOXIDE 

50.  Carbon,  a.  Arrange  an  apparatus  as  shown  in  Fig.  28. 
Partly  fill  the  hard-glass  test  tube  A  with  small  pieces 
of  wood  —  preferably  with  hardwood  sawdust.  Heat  the 
wood,  gently  at  first  and  then 
more  strongly.  After  some 
liquid  has  collected  in  B, 
apply  a  light  to  the  jet  C. 
Are  combustible  gases  evolved 
on  distilling  wood  ?  Continue 
the  heating  as  long  as  gas  is 
evolved,  then  examine  the 
liquid  collecting  in  B.  Is  it 
soluble  in  water?  How  would 
you  describe  its  odor  ?  What 
is  the  residue  left  in  A  ?  Will 
it  burn  ?  Does  it  leave  any 
ash? 

&.  Bring    a    cold    porcelain 
dish    into    a   small,    luminous 
Bunsen    flame.    Note    the    de- 
posit.   What  is  this  form  of  carbon  called  ?   In  what  other 
forms  does  carbon   exist  ?    What   properties  have  all  these 
forms  in  common  ? 

c.  Put  one  fourth  of  a  test  tube  full  of  bone  black  into  a 
small  flask  and  pour  over  it  about  50  cc.  of  water,  to  which 
has  been  added  a  few  drops  of  a  solution  of  litmus  or  of 
indigo.  Thoroughly  mix  the  contents  of  the  flask,  then  heat 
it  gently  for  a  few  minutes,  and  filter.  If  the  filtrate  is  not 

[36] 


FIG.  28 


J 


V-v\ 


jU^X  ^   03  .  ' 


jU^.^  fflr^Lf.^ 


decolorized,  repeat,  using  more  bone  black.  What  is  the  com- 
position of  bone  black  ?  By  what  other  name  is  it  known  ? 
What  use  does  this  experiment  suggest  for  it? 

d.  Is  carbon  at  ordinary  temperatures  an  active  element  ? 
Test  it  with  the  common  acids.    How  does  the  charring  of 
wood  preserve  it  ? 

e.  Prepare  an  apparatus  according  to  Fig.  29.  The  bottle  A 
contains  a  solution  of  sodium  hydroxide,  while  B  and  C  con- 
tain a  solution  of  calcium  hydroxide  (lime water).    The  hard- 
glass  tube  D  contains  one  or  two  small  pieces  of  charcoal. 
Charge  your  oxygen  generator  as  in  §  15  and  connect  the 
rubber  delivery  tube  with  E.    Now  pass  a  slow  current  of 
oxygen  through  the  apparatus,  at  the  same  time  heating  the 
charcoal   in  D  until  it  just  begins  to  glow.     Describe  the 

results.     What    do    they 

V  D 

prove  ?     Why    pass    the  — 

oxygen  through  the  solu- 
tions in  A  and  B  ?  What 
reactions  take  place  in  D 
andC(R)?  T"  "B  C 

f.  In  a  hard-glass  test  FIG.  29 
tube  heat  an  intimate  mix- 
ture of  2  or  3  g.  of  black  copper  oxide  and  an  equal  bulk  of 
powdered  charcoal.    Pass  the  evolved  gases  through  a  little 
limewater  in  a  test  tube.    Write  all  the  equations  involved 
in  the  reactions.    What  remains  in  the  test  tube  ?    How  can 
you  prove  it  ?    What  use  does  this  suggest  for  charcoal  ? 

51.  Carbon  dioxide,  a.  Put  some  pieces  of  marble  in  your 
hydrogen  generator,  cover  them  with  water,  and  add  a  little 
concentrated  hydrochloric  acid  through  the  funnel  tube.  Fill 
two  or  three  bottles  with  the  gas  evolved  by  downward  dis- 
placement. (To  test  whether  the  bottles  are  filled  or  not, 
hold  a  burning  splint  in  the  mouth  of  the  bottle.) 

b.  Devise  an  experiment  to  show  whether  the  gas  is  heavier 
or  lighter  than  air.  Attempt  to  pour  it  from  one  bottle  to 
another,  as  you  would  a  liquid,  and  test  with  a  burning 
splint  for  its  presence  in  the  second  bottle. 

[37] 


*^T  xUT 


c.  Half  fill  a  small  beaker  with  limewater  and  pass  carbon 
dioxide  through  the  liquid  (R).   The  formation  of  a  white  pre- 
cipitate in  limewater  is  a  test  for  carbon  dioxide. 

d.  By  blowing  through  limewater  prove  the  presence   of 
carbon  dioxide  in  the  air  exhaled  from  the  lungs  (R).   Hold  a 
wide-mouthed  bottle  above  a  small  flame  so  that  the  hot  gase- 
ous products  of  combustion  will  collect  in  it ;  then  quickly  add 
a  few  cubic  centimeters  of  limewater,  cover  the  mouth  of  the 
bottle  with  the  hand,  and  shake  up  the  contents.    What  do 
the  results  prove  ?  What  are  the  sources  of  carbon  dioxide 
in  the  air  ? 

52.  Weight  of  i  liter  of  carbon  dioxide.  Obtain  a  250-cc. 
Erlenmeyer  flask  (Fig.  30)  and  a  well-rolled  cork  to  fit  it 
tightly.  Insert  a  small  glass  tube  through  the  cork.  A  piece 
of  pure  rubber  tubing  is  slipped  over  the 
end  of  the  glass  tubing  and  the  rubber 
tubing  is  closed  with  a  pinch  clamp  as  shown 
in  the  figure.  Thoroughly  dry  the  flask 
both  inside  and  out,  then  insert  the  stopper 
tightly  and  mark  the  point  to  which  it  ex- 
tends in  the  flask  by  a  strip  of  paper  cut 
from  a  label.  Open  the  clamp  for  a  moment 
to  adjust  the  pressure  of  the  air.  Accurately  FlG  "30 

weigh  to  third  decimal  place  the  flask 
and  stopper,  recording  the  reading  of  the  barometer  and  the 
temperature  of  the  air.  Ne^t  arrange  an  apparatus  as  shown 
in  Fig.  31.  A  represents  the  carbon  dioxide  generator,  B  repre- 
sents a  bottle  partially  filled  with  sulfuric  acid,  and  C  repre- 
sents the  250-cc.  flask  (cork  removed).  Now  generate  a  slow 
current  of  carbon  dioxide  in  A.  The  gas  is  dried  as  it  bubbles 
through  the  sulfuric  sa/Sjjf^n  B  and  is  then  conducted  into  (7, 
gradually  displacing  the  air  in  C.  Continue  passing  the  gas 
into  (7,  until  you  are  satisfied  that  all  the  air  has  been  expelled ; 
then  while  the  gas  is  still  being  evolved,  slowly  remove  the 
tube  D  from  the  flask,  and  at  once  insert  the  cork  in  the  flask 
to  the  point  indicated  by  the  gummed  paper.  Open  the  clip 
for  a  moment  to  adjust  the  pressure ;  then  weigh  the  flask  as 

[38] 


**Jr~~ 

*^      l£&?T. 


before.  Next  completely  fill  the  flask  with  water  and  insert 
the  cork,  allowing  the  extra  water  to  run  out  through  the 
rubber  tubing ;  then  measure  the  water  by  transferring  it  to 


FIG.  31 

a  graduated  cylinder.  This  will  give  you  the  volume  of  the 
flask.  Insert  the  appropriate  values  in  the  table  below  and 
calculate  the  weight  of  1  liter  of  carbon  dioxide. 

i    s  i  ^  r 

1.  Barometric  pressure / 

2.  Temperature ^  I 

3.  Volume  of  flask 

4.  Weight  of  flask  filled  with  air  under  laboratory  conditions  3  ^  »  5 

5.  Weight  of  1  liter  of  air  under  standard  conditions       .     .  1.293  g. 


6.  Weight  of  1  liter  of  air  under  laboratory  conditions     .     .    / 

7.  Weight  of  flask  empty.    (Calculate  from  3,  5,  and  6)  . .     .3 

8.  Weight  of  flask  filled  with  carbon  dioxide  under  labo- 

ratory conditions 

9.  Weight  of  carbon  dioxide  in  flask 

10.   Weight  of  1  liter  of  carbon  dioxide  under  standard  condi- 
tions.   (Calculate  from  3  and  9) 


'    t 


[39] 


CHAPTER  VIII 
NITROGEN  AND  THE  ATMOSPHERE 

53.  The  preparation  and  properties  of  nitrogen.  (On  account 
of  its  great  affinity  for  oxygen,  phosphorus  must  be  kept  and 
handled  under  water.  Never  bring  the  dry  substance  in  con- 
tact with  the  skin,  as  it  may  ignite  and  cause  a  serious  burn.) 
a.  Cover  the  bottom  of  a  pneumatic  trough  with  water  to  a 
depth  of  2  or  3  cm.  Float  on  the  water  a  porcelain  crucible 
containing  a  small  piece  of  phosphorus.  Ignite  the  phos- 
phorus by  touching  it  with  the  hot  end  of  a  wire  or  file, 
and  quickly  invert  over  the  crucible  a  large  beaker  or  wide- 
mouthed  bottle,  being  careful  to  keep  the  rim  of  the  beaker 
below  the  surface  of  the  water.  The  white  fumes  formed  have 
the  composition  P2O&.  Leave  the  beaker  in  position  until  the 
fumes  have  entirely  disappeared.  Note  that  the  water  has 
risen  in  the  beaker.  Explain.  Adjust  the  beaker  or  the  water 
in  the  trough  so  that  the  level  of  the  liquid  inside  and  out- 
side of  the  beaker  is  the  same ;  then  cover  the  beaker  with  a 
glass  plate  and  turn  it  into  an  upright  position.  Test  the  gas 
with  a  burning  splint. 

b.  In  a  250-cc.  flask  place  a  mixture  of  3  g.  of  ammonium 
chloride  and  6  g.  of  sodium  nitrite,  and  add  20  cc.  of  water. 
Provide  the  flask  with  a  cork  (one-hole)  and  delivery  tube, 
so  that  the  gas  evolved  may  be  collected  over  water  as  in  the 
case  of  oxygen  and  hydrogen.  Have  at  hand  a  vessel  of  cold 
water  so  that  the  flask  may  be  cooled  by  lowering  it  into  the 
water  in  case  the  action  becomes  too  violent.  Clamp  the  flask 
on  a  ring  stand  and  apply  a  very  gentle  heat,  moving  the 
burner  about  with  the  hand.  As  soon  as  the  action  begins, 
withdraw  the  burner.  After  the  air  has  been  expelled  from 

[40] 


/ 

^^  , ^  .,    ,    ^^^t-^s^»^^^ 

^c^^  #**4r^  ^^^^-^-  ~^ 

v^X^^—  r  4^-x— 


the  apparatus  fill  two  or  three  bottles  (250-cc.)  with  the  gas. 
If  the  action  becomes  too  violent,  immerse  the  flask  in  cold 
water.  The  reaction  which  takes  place  is  expressed  in  the 
following  equation: 

NH4C1  +  NaNO2  =  NH4NO2  +  NaCl 

The  ammonium  nitrite  then  decomposes  into  water  and  nitro- 
gen (R).  (The  symbol  (R)  indicates  that  the  equation  for 
the  reaction  is  to  be  written.)  Note  the  physical  properties 
of  the  gas.  Test  with  a  burning  splint. 

54.  Determination  of  the  relative  volumes  of  nitrogen  and 
of  oxygen  in  the  air.1  This  determination  may  be  made  by 
bringing  in  contact  with  a  definite  volume  of  air  a  liquid 
which  not  only  absorbs  the  oxygen  but 
in  doing  so  flows  into  the  tube  which 
contains  the  air  and  fills  a  space  equal 
to  that  previously  occupied  by  the  oxy- 
gen. The  volume  of  this  liquid  can  be 
easily  measured,  and  in  this  way  the 
volume  of  the  absorbed  oxygen  may 
be  ascertained. 

The  solution  used  to  absorb  the  oxy- 
gen soon  loses  its  strength  on  exposure 
to  the  air  ;  hence  the  experiment  must 
be  performed  rapidly.  Before  preparing 
the  solution  the  student  should  prac- 
tice the  manipulations  involved  in  the 
experiment. 

Prepare  an  apparatus  according  to  Fig.  32.  In  this  apparatus 
C  represents  a  test  tube  about  15  cm.  in  length  (use  the  hard- 
glass  tube  employed  in  the  preparation  of  oxygen).  The  tube 
is  fitted  with  a  two-hole  rubber  stopper.  Through  one  of  the 
holes  is  fitted  a  glass  tube,  which,  together  with  the  accom- 
panying rubber  tube  and  funnel,  is  the  same  as  was  used  to 
introduce  the  acid  into  the  bottle  A  (Fig.  27).  The  remain- 
ing hole  in  the  stopper  is  closed  with  a  glass  rod.  Notice  that 

1  From  Cooley's  "Laboratory  Studies." 
[41] 


FIG.  32 


both  the  glass  rod  and  the  tube  extend  just  through  the  small 
end  of  the  stopper.  Now  close  the  rubber  tube  tightly  with  a 
screw  clamp  B.  Disconnect  the  test  tube  and  remove  the  glass 
rod  from  the  stopper,  preparatory  to  performing  the  experi- 
ment. Prepare  an*  alkaline  solution  of  pyrogallic  acid,  as 
follows :  Dissolve  5  g.  of  potassium  hydroxide  in  5  cc.  of 
water  and  cool  the  solution  to  room  temperature.  Add  this 
to  a  solution  of  4  g.  of  pyrogallic  acid  in  10  cc.  of  water,  and 
at  once  pour  the  resulting  liquid  into  the  funnel  A.  Quickly 
open  the  screw  clamp  until  both  the  rubber  tube  and  the 
glass  tube  are  filled  with  the  liquid,  then  close  tightly.  Con- 
nect the  test  tube,  holding  it  by  the  rim  to  avoid  heating  the 
contained  air,  and  insert  the  glass  rod  in  the  cork.  The  air 
inclosed  in  the  tube  is  now  at  the  same  temperature  and 
pressure  as  the  surrounding  air.  Now  open  the  screw  clamp. 
The  liquid  flows  in,  absorbing  the  oxygen.  When  the  liquid 
ceases  to  enter,  grasp  the  tube  by  the  rim  and  invert  it  as 
shown  in  the  dotted  lines  of  the  figure,  adjusting  it  so  that 
the  level  of  the  liquid  is  the  same  in  both  the  tube  and  the 
funnel  (why?).  Then  clamp  the  rubber  tube  and  return  the 
test  tube  to  its  original  position.  Mark  the  volume  of  the  air 
originally  inclosed  in  «the  tube  by  placing  a  narrow  strip  of 
gummed  paper  about  the  tube  at  the  lower  end  of  the  stop- 
per ;  also  mark  by  a  strip  of  paper  placed  at  the  level  of  the 
liquid  in  the  tube,  the  volume  of  the  oxygen  absorbed.  Dis- 
connect the  tube  and  rinse  it.  Measure  the  volume  of  the  tube 
to  each  strip  of  paper  by  pouring  in  water  from  a  graduated 
cylinder.  From  these  data  calculate  the  volume  of  oxygen 
and  of  nitrogen  in  100  volumes  of  air. 

NOTE.  This  experiment  disregards  the  presence  in  air  of  all  constitu- 
ents other  than  oxygen  and  nitrogen.  The  volume  of  such  constituents, 
however,  in  the  volume  of  air  taken,  is  so  small  that  it  may  be  neglected. 

55.  Other  constituents  of  the  air.  a.  Expose  a  piece  of  cal- 
cium chloride  to  the  air  for  two  or  more  hours.  Explain. 

b.  Expose  a  few  cubic  centimeters  of  limewater  to  the  air 
for  a  half  hour.  Explain  the  results. 

[42] 


CHAPTER  IX 
SOLUTIONS 

56.  Rate  of  solution,    a.  Drop  a  crystal  of  potassium  per- 
manganate  into    a   test  tube  partly  filled  with  water  and 
shake  the  liquid  until  the  solid  is  dissolved.    Note  the  color 
of  the  solution. 

b.  Place  a  test  tube  filled  with  water  in  a  test-tube  rack 
and  drop  into  it  a  crystal  of  potassium  permanganate.  Allow 
the  liquid  to  remain  (without  shaking  it)  until  near  the  close 
of  the  laboratory  period.  Note  the  result.  What  does  the 
experiment__show  ? 

57.  Solubility,    a.  To  a  test  tube  two  thirds  full  of  boiling 
water  add  common  salt,  small  portions  at  a  time,  as  long  as 
any  dissolves.    If  an  excess  remains  undissolved  after  boiling 
a  short  time,  add  just  enough  water  to  bring  it  into  solution. 
Immerse  the  test  tube  in  a  beaker  of  cold  water.   Is  salt  more 
soluble  in  cold  or  in  hot  water  ?  Repeat  the  experiment,  using 
powdered  potassium  chloride  or  potassium  nitrate  instead  of 
salt,  immersing  the  test  tube  in  the  same  beaker  as  before. 
When  the  two  tubes  have  cooled,  note  the  relative  quantities  of 
the  solids  that  have  separated.   What  inference  do  you  draw  ? 

b.  Filter  the  salt  solution  through  a  dry  filter  into  a  dry 
test  tube  and  note  the  temperature.    Weigh  a  small  evaporat- 
ing-dish  and  watch-glass  cover;  then  pour  10  or  15  cc.  of  the 
clear  salt  solution  into  the  dish  and  reweigh.  Now  evaporate  to 
dryness,  taking  note  of  the  precautions  given  in  §  46.    When 
the  dish  is  cool,  reweigh.    From  your  results  calculate  the 
solubility  of  common  salt  at  the  indicated  temperature. 

c.  In  a  similar  way  determine  the  sol  ability  of  the  potas- 
sium nitrate  or  of  potassium  dichromate. 

[43] 


4"#^# 


^JL^Ui- 


58.  Supersaturated  solutions.  To  20  cc.  of  water  in  a  beaker 
add  20  g.  of  hydrated  sodium  sulfate  and  gently  warm  until 
solution  is  complete.    Consult  the  solubility  curve  of  sodium 
sulfate.     How  high  is  it  advisable  to  heat  the  solution  ?    Is 
the  solution  mobile  or  viscous  ?    Pour  the  solution  into  two 
test  tubes  and  set  it  aside  to  cool.  If  no  crystallization  occurs 
at  room  temperature,  the  solution  should  be  supersaturated. 
Add  a  very  small  crystal  of  the  solid  sodium  sulfate.    Why 
do  crystals  form  ?    In  what  condition  is  the  remaining  solu- 
tion ?     Define   supersaturation.    A  concentrated  solution  of 
ferric  nitrate  in  dilute  nitric  acid  saturated  at  about  60°  is 
more  certain  than  the  hydrated  sodium  sulfate  to  work  well. 

59.  Solubility  of  liquids.    Pour  about  5  cc.  of  water  into  a 
test  tube,  add  an  equal  volume  of  kerosene,  and  mix  the 
liquids  by  shaking  the  test  tube.    Set  the  tube  aside  for  a 
few  minutes   and  note  the  result.    Repeat  the  experiment, 
substituting  alcohol  for  kerosene.    Note  the  result. 

60.  Fractional  distillation.    Distill  a  mixture  of  10  cc.  of 
alcohol  (boiling  point  78.3°)  and  30  cc.  of  water.    Collect  the 
first  1  or  2  cc.  of  the  distillate  in  an  evaporating-dish  and  test 
with  a  flame.    In  the  same  way  test  successive  portions  of 
the  distillate.    Does  there  seem  to  be  a  separation  of  the  two 
liquids  ?    By  the  above  method  a  mixture  of  liquids  boiling 
at  different  temperatures  may  often  be  separated  more  or  less 
perfectly.    What  name  is  given  to  this  process  ? 

61.  Constant-boiling  solutions.  To  1  or  2  cc.  of  dilute  hydro- 
chloric acid  add  a  drop  or  two  of  a  solution  of  silver  nitrate.  In 
this  case  the  formation  of  a  white  precipitate  (silver  chloride) 
shows  the  presence  of  hydrochloric  acid  in  the  solution.  Pour 
about  75  cc.  of  dilute  hydrochloric  acid  into  a  distilling-flask 
connected  with  a  Liebig  condenser  and  slowly  distill,  catching 
the  distillate  in  a  test  tube.    When  about  5  cc.  has  been  col- 
lected, test  it  with  1  or  2  drops  of  silver  nitrate  solution.    Is 
there  an  appreciable  precipitate  ?    Continue  the  distillation, 
testing  each  portion  of  5  cc.  as  it  is  collected.  Does  the  quan- 
tity of  the  precipitate  formed  increase  gradually  or  suddenly? 
When  a  decided  precipitate  is  obtained,  collect  about  2  cc.  of 

[44] 


the  distillate  and  set  it  aside ;  then,  without  further  testing, 
distill  over  as  much  of  the  solution  as  you  safely  can,  collect- 
ing the  last  2  cc.  See  that  these  two  portions  are  as  nearly 
equal  as  possible,  then  dilute  each  to  about  15  cc.,  and  to 
5  cc.  of  each  add  silver  nitrate  as  long  as  a  precipitate  forms. 
Allow  the  precipitate  to  settle,  and  compare  the  two  quan- 
tities. From  all  of  your  observations  do  you  think  it  would 
be  possible  to  fraction  the  more  concentrated  solutions  of 
hydrochloric  acid? 


[45] 


CHAPTER  X 


CHLORINE ;  HYDROGEN  CHLORIDE ;    HYDROCHLORIC  ACID 

„.  62.  Preparation  of  chlorine.  (All  the  following  experiments 
must  be  performed  in  the  hood,  and  great  care  taken  not  to  inhale 
the  gas.)  a.  Place  about  1  g.  of  manganese  dioxide  in  a  test 
tube,  add  2  or  3  cc.  of  hydrochloric  acid,  and  heat  gently  (R). 

b.  In  a  similar  way  test  the  action  of  hydrochloric  acid 
upon  lead  peroxide,  potassium  permanganate,  and  potassium 
dichromate.    Will  all  compounds 

containing  oxygen  oxidize  hydro- 
chloric acid  ?  Try  lead  oxide 
(litharge)  and  sodium  sulfate. 

c.  Repeat  experiment  a,  replac- 
ing the  hydrochloric  acid  with  small 
amounts  of  common  salt  and  sul- 
furic  acid.    The  manganese  dioxide 
and  salt  should  first  be  mixed  to- 
gether and  the  sulfuric  acid  added 
to  the  mixture.    Compare  the  reac- 
tions in  a  and  c  (R).    Which  of 
the  two  methods  is  the  cheaper  for 
preparing  chlorine  ?    Give  reason 
for  your  answer. 

d.  Usual  laboratory  method  (two 

students  working  together).  Arrange  an  apparatus  according 
to  Fig.  33,  using  a  250-cc.  flask.  Thoroughly  clean  and  dry 
five  wide-mouthed  bottles  (250-cc.)  for  collecting  the  gas. 
Place  in  the  flask  from  20  to  25  g.  of  manganese  dioxide. 
Insert  the  cork  and  pour  150  cc.  of  hydrochloric  acid  through 
the  funnel  tube.  Shake  the  flask  so  as  to  thoroughly  wet  the 

[46] 


FIG.  33 


0 

— : > 


xH 


'Ao^*  ^^M^-^-^*^*-^-   ^C~^l  tfjti/^*  o^-rte^  fa^ 

•^H^JMA^ 


manganese  dioxide.  Warm  gently,  applying  just  enough  heat 
to  cause  a  gentle  evolution  of  the  gas  but  not  enough  to  boil 
the  liquid.  Fill  all  the  bottles  with  the  gas,  collecting  it  by 
downward  displacement;  then  prepare  some  chlorine  water 
by  bringing  the  exit  tube  into  a  bottle  containing  a  little 
water,  so  that  the  gas  bubbles  up  through  the  liquid.  Con- 
tinue the  heating  until  no  more  chlorine  is  evolved  (chlorine 
water). 

'  e.  A  part  of  the  class  should  prepare  chlorine  as  follows: 
Instead  of  the  funnel  tube  (Fig.  33)  use  a  dropping-funnel 
(an  ordinary  funnel  connected  with  a  straight  glass  tube  by 
a  piece  of  rubber  tubing  may  be  used,  a  screw  clamp  on  the 
rubber  connection  taking  the  place  of  the  stopcock,  as  shown 
in  Fig.  17).  Place  about  20  g.  of  potassium  permanganate 
crystals  in  the  dry  flask  and  slowly  admit  concentrated  hydro- 
chloric acid  from  the  dropping-funnel,  regulating  the  flow  of 
the  acid  by  the  rate  at  which  chlorine  is  evolved. 

63.  Properties  of  chlorine,  a.  Sprinkle  a  pinch  of  finely 
powdered  antimony  into  one  of  the  bottles  of  the  gas.  SbCl3 
is  formed.  Is  this  reaction  an  example  of  combustion  ?  Is 
the  presence  of  oxygen  necessary  for  combustion  ? 

b.  Support  by  forceps  a  small  piece  of  copper  foil,  heat  it 
to  redness,  and  immediately  drop  it  into  a  bottle  of  the  gas. 
Describe  the  result.    What  is  formed? 

c.  In  a  bottle  of  the  dry  gas  suspend  a  strip  of  colored 
calico  and  two  strips  of  paper,  one  with  dry  writing  in  ink 
on  it,  the  other  with  printing  (printer's  ink). 

d.  Repeat  c,  using  similar  strips  moistened  with  water. 
Describe  the  results  in  c  and  in  <L    What  part  does  the  water 
play  in  the  bleaching? 

e.  Test  the  action  of  the  gas  on  strips  of  red  and  of  blue  lit- 
mus paper.    Generate  hydrogen  by  the  usual  method.    After 
taking  the  required  precautions  (what  are  they  ?)  ignite  the 
hydrogen  delivered  from  a  jet  and  introduce  the  flame  into 
a  bottle  of  chlorine.    Does  the  hydrogen  continue  to  burn  ? 
What  is  formed  ?    Allow  the  hydrogen  to  burn  until  the 
chlorine  is  all  used  up ;  then  test  with  moist  litmus  paper(?). 

[47] 


I '3 


,  , 

/1,,  /r- 

^ 


J, 

'**£&& 

/(^rtAe^fc-   ;l 

A 

/XVV^A^xO'       '    AA      'JxO  Z^. 

( 

/ 

J  /x  ./     >.^  yCc^' 

^>t,..,        ^c^  *X^tf£^^ 


64.  Preparation  of  hydrogen  chloride  and  hydrochloric  acid. 
Connect  the  generator  used  in  the  preparation  of  chlorine 
(Fig.  33)  with  a  wide-mouthed  bottle,  as  shown  in  Fig.  34. 
The  delivery  tube  F  is  cut  at  A  and  at  Z>,  and  the  ends 
(rounded)  are  joined  by  rubber  tubing.   Pour  about  30  cc.  of 
water  into  the  bottle  B.    Notice  that  the  tube  extending  into 
the   bottle  does  not  quite  touch 

the  surface  of  the  water.  Pre- 
pare some  dilute  sulfuric  acid 
by  carefully  pouring  25  cc.  of 
the  concentrated  acid  into  10  cc. 
of  water  (note  precaution,  §  23). 
The  mixture  must  be  stirred 
and  cooled  from  time  to  time 
while  the  acid  is  being  added. 
Have  at  hand  two  dry  bottles 
in  which  to  collect  some  of  the 
gaseous  hydrogen  chloride.  Put 
about  50  g.  of  common  salt  into 
the  generator  flask,  insert  the 
cork,  pour  the  cold  dilute  sul- 
furic acid  through  the  funnel  tube,  mix  the  contents  by  a 
gentle  motion  of  the  flask,  and  after  two  or  three  minutes 
warm  gently  with  a  small  flame.  Notice  the  currents  in  the 
water  in  B.  What  is  the  cause  of  them  ?  As  soon  as  the  gas 
is  evolved  regularly,  disconnect  the  generator  flask  at  D  long 
enough  to  collect  (by  downward  displacement)  two  bottles  of 
the  gas.  Cover  these  tightly  with  glass  plates  and  set  aside ; 
then  again  connect  the  generator  with  B  and  continue  to 
apply  a  gentle  heat  as  long  as  any  gas  is  evolved.  What 
two  reactions  are  possible?  From  the  weights  of  materials 
used  and  the  conditions  of  the  experiment,  which  one  would 
you  expect  to  take  place?  Why  not  extend  the  tube  in  B 
to  the  bottom  of  the  bottle  ? 

65.  Properties  of  hydrogen  chloride,    a.  What  is  the  color 
of  gaseous  hydrogen  chloride.    Account  for  the  formation  of 
fumes  when  the  gas  is  exposed  to  the  air. 

[48] 


FIG.  34 


X  /t  tl^  <' 


3^^v^v>^^ 


uw 

±7 

. 


b.  Test  with  a  lighted  splint   one  of  the  bottles  of  gas 
collected  in  the  experiment  of  §  64.    Is  it  combustible  ?    Is 
it  a  supporter  of  combustion  ? 

c.  Fill  a  large  beaker  with  water  and  color  it  with  a  few 
drops  of  a  solution  of  blue  litmus.    Now  uncover  the  remain- 
ing bottle  of  gas  collected  in  the  experiment  of  §  64,  invert 
it,  and  at  once  bring  the  mouth  of  it  under  the  surface  of 
the  water  in  the  beaker.    Describe  the  results.    What  does 
the  experiment  prove  ? 

66.  Properties  of  hydrochloric  acid.  a.  Put  a  drop  of  the 
aqueous  solution  of  the  acid  from  bottle  B  (§  64)  on  a  bit  of 
blue  litmus  paper.  Note  the  result.  Add  two  drops  of  the 
solution  to  3  or  4  cc.  of  water  and  taste  a  drop.  Make  a 
test-tube  experiment  to  prove  the  presence  of  chlorine  in  the 
acid ;  also  one  to  prove  the  presence  of  hydrogen.  Write  the 
equations  in  each  case.  How  does  the  solution  compare  with 
the  hydrochloric  acid  on  your  desk  ? 

b.  Pour  out  some  of  the  solution  from  the  generator  flask 
into  a  beaker  and  allow  it  to  cool.    Add  some  concentrated 
hydrochloric   acid.     What    change    do    you   note  ?     Is    this 
reaction  reversible  ? 

c.  For  the  determination  of  the  composition  of  hydrogen 
chloride,  see  §  113. 


[49] 


ux^ 


xi 


, 


-^^ 

^V 


CHAPTER  XI 
SODIUM;  SODIUM  HYDROXIDE 

67.  Sodium ;  sodium  hydroxide,  a.  Recall  §  20.  Obtain  a 
small  piece  of  sodium  from  the  instructor.  Cut  it  and  note 
the  rapidity  with  which  the  freshly  cut  surface  is  tarnished. 
Half  fill  an  evaporating-dish  with  water,  then  drop  the  sodium 
into  it  and  quickly  cover  the  dish  with  a  glass  plate.  After 
the  action  (R)  has  entirely  ceased,  test  the  liquid  with  red 
litmus  paper.  What  compound  is  dissolved  in  the  water  ? 

b.  Weigh  out  about  5  g.  of  good  quicklime  and  moisten  it 
with  about  an  equal  weight  of  water  in  an  evaporating-dish, 
warming  it,  if  necessary,  to  make  it  slake  (R).  When  it  has 
ceased  slaking,  dilute  to  100  cc.  and  add  the  required  weight 
of  sodium  carbonate  to  produce  complete  double  decomposi- 
tion (R).  Boil  for  a  few  minutes  to  render  filtration  easier, 
and  then  filter.  What  is  the  filtrate  ?  How  is  this  obtained 
commercially  in  solid  form  ?  Test  a  little  of  the  filtrate  to  see 
if  an  excess  of  carbonate  has  been  added  (how  will  you  do 
this  ?),  and  if  an  excess  is  found,  devise  a  way  to  remove  it. 
Add  2  drops  of  the  filtrate  to  5  cc.  of  water  and  taste  the 
solution  (?).  Dip  a  piece  of  red  litmus  paper  into  the  filtrate  (?). 


[50] 


A4- 


0    - 


jP 


• 


CHAPTER  XII 
ACIDS;   BASES;   SALTS;   NEUTRALIZATION 

68.  Acids,    a.  Recall  the  properties  of  hydrochloric  acid 
(§  66).    Prepare  dilute  solutions  of  each  of  the  following 
acids  by  adding  2  cc.  of  each  of  the  concentrated  acids  to 
10  cc.  of  water  and  thoroughly  mixing:  hydrochloric,  nitric, 
sulfuric,  acetic.    By  means  of  a  clean  glass  rod  transfer  a 
drop  of  each  to  a  piece  of  blue  litmus  paper  and  then  to  a 
piece  of  red.    What  changes  do  you  notice? 

b.  Recall  the  action  of  hydrochloric,  sulfuric,  and  nitric 
acids  on  zinc  (§§  22,  23).    Nitric  acid  is  a  strong  oxidizing 
agent.    Might  this  account  for  the  fact  that  hydrogen  is  not 
evolved  when  this  acid  reacts  with  zinc  ? 

c.  Add  one  drop  of  hydrochloric  acid  solution  to  10  cc.  of 
water.    Stir  thoroughly  and  taste  the  solution.    Repeat,  using 
acetic  acid  (?).    Compare  the  formulas  of  the  acids.    In  what 
respect  are  the  acids  similar  in  composition? 

69.  Bases.    Prepare  dilute  solutions  of  each  of  the  follow- 
ing bases:  sodium  hydroxide,  potassium  hydroxide,  calcium 
hydroxide.    Try  the  effect  of  each  of  these  solutions  on  blue 
and  on  red  litmus  paper.    Taste  a  drop  of  the  calcium  hydrox- 
ide solution.    Compare  the  formulas  of  the  above  bases.    In 
what  respects  are  the  bases  similar  in  composition  ?    From 
your  results  characterize  the  properties  of  bases  on  the  sup- 
position that  the  above  compounds  are  typical  bases. 

70.  Salts.    Dilute  5  cc.  of  the  ordinary  laboratory  solution 
of  sodium  hydroxide  (1  part  of  the  solid  to  10  parts  of  water) 
with  an  equal  volume  of  water.    To  this  solution  add  a  few 
drops  of  hydrochloric  acid;  the  two  react  with  evolution  of 
heat  (R).     Stir  the  resulting  solution  with  a  glass  rod  and 
test  its  action  on  blue  and  on  red  litmus  paper.    Has  it  acid 

[51] 


or  basic  properties  ?  Continue  to  add  the  acid,  drop  by  drop, 
with  stirring,  till  the  resulting  solution  is  neutral  or,  at  most, 
slightly  acid.  Pour  the  solution  into  an  evaporating-dish  and 
evaporate  to  dryness.  What  compound  remains  ?  What  is 
the  name  given  to  the  compounds  formed  by  the  action  of 
acids  with  bases  ?  Characterize  these  compounds. 

71.  Ratio  of  acid  to  base  in  neutralization.  Prepare  a  dilute 
solution  of  sodium  hydroxide  by  diluting  20  cc.  of  the  lab- 
oratory reagent  to  100  cc. ;  also  a  dilute  solution  of  sulfuric 
acid  by  adding  1  cc.  of  the  concentrated  acid  to  100  cc.  of 
water.  Rinse  out  a  burette  (storeroom),  first  with  distilled 
water,  then  with  a  little  of  the  alkaline 
solution.  Support  the  burette  (Fig.  35) 
and  pour  the  alkaline  solution  into  it 
until  the  level  of  the  liquid  is  2  or 
3  cm.  above  the  zero  mark.  Turn  the 
stopcock  and  let  the  solution  flow  out 
until  the  bottom  of  the  curved  surface 
(meniscus)  is  on  a  level  with  the  zero 
mark.  In  a  similar  way  fill  a  second 
burette  with  the  acid  solution.  Now 
let  exactly  15  cc.  of  the  acid  solution 
flow  into  a  small  beaker,  add  2  drops 
of  a  solution  of  phenolphthalein  and 
run  in  2  or  3  cc.  of  the  alkaline  solu- 
tion. Notice  that  where  the  liquids 
come  in  contact,  a  reddish  color  is 
produced,  which  disappears  quickly 
on  stirring.  Run  in  more  of  the  solu- 
tion, a  little  at  a  time,  until  the  color  fades  slowly,  then  a 
drop  at  a  time  until  the  entire  liquid,  on  stirring,  remains 
colored  faintly  red.  This  marks  approximately  the  point  of 
neutralization.  Note  the  number  of  cubic  centimeters  of  the 
alkaline  solution  used.  Repeat  the  experiment,  using  different 
volumes  of  acid,  say  10  cc.  and  20  cc.  Calculate  in  each  case 
the  number  of  cc.  of  the  alkaline  solution  required  to  neutralize 
1  cc.  of  the  acid  solution.  What  do  the  results  prove  ? 

[52] 


FIG.  35 


CHAPTER  XIII 
IONIZATION 

72.  Conducting  power  of  various  solutions,  a.  Obtain  a 
conductivity  apparatus  (Fig.  36)  from  the  storeroom  and  pol- 
ish the  copper-wire  electrodes  with  emery  paper  until  they 
are  bright  and  free  from  oxide.  At  the  beginning  of  each 
experiment  see  that  the  electrodes  are  bright  and  dry,  and 


FIG. 36 

that  the  cell  A  is  also  perfectly  clean  and  dry.  Unscrew  a 
lamp  C  from  a  convenient  socket  above  your  desk,  screw  it 
loosely  into  the  socket  on  your  apparatus,  and  attach  the 
apparatus  to  the  empty  socket  on  the  lighting  system  by  means 
of  the  extension  cord  and  plug  B.  Every  time  a  change  is  to 
be  made  in  the  cell,  loosen  the  lamp  in  the  socket,  and  do  not 
screw  it  down  to  make  contact  until  all  the  connections  of 
the  cell  have  been  arranged. 

b.  Partly  fill  the  cell  A  with  dry  powdered  salt,  dip  the 
electrodes  into  the  powder,  arrange  the  connections  at  the 

[53] 


binding-posts,  and  screw  down  the  lamp  C.    Have  you  any 
evidence  that  the  powdered  salt  is  a  conductor  ? 

c.  In  a  similar  way  test  the  conductivity  of  distilled  water 
and  tap  water. 

d.  Dissolve  about  1  g.  of  the  salt  in  10  cc.  of  distilled  water 
and  test  the  conductivity  of  the  solution.    Account  for  the 
result  as  compared  with  the  results  obtained  in  b  and  c. 

e.  Repeat  b  and  d,  substituting  sugar  for  salt.    Account  for 
any  difference  in  results. 

73.  Acids  from   the   standpoint  of   the    ionization   theory. 
a.  Test  the  conductivity  of  dilute  solutions  of  the  following 
acids:  hydrochloric,  sulfuric,  acetic. 

b.  Ask  the  instructor  for  a  solution  of  dry  hydrogen 
chloride  in  benzene  or  toluene.  Determine  whether  it  con- 
ducts the  electric  current  (care  must  be  taken  to  exclude 
all  moisture).  Pour  a  little  of  the  solution  oil  a  clean,  dry 
iron  nail.  Repeat,  using  the  ordinary  aqueous  solution.  Ac- 
count for  the  difference  in  action  between  the  two  solutions. 
Characterize  acids  from  the  standpoint  of  the  ionization  theory 
on  the  basis  that  the  above  acids  are  typical  ones  and  that  all 
others  act  like  them. 

74.  Relative  strength  of  acids,    a.  Obtain  from  your  in- 
structor about  15  cc.  of  twice  normal  solutions  of  each  of  the 
following  acids :  hydrochloric,  sulfuric,  and  acetic.    Test  their 
relative  conductivities  (Fig.  36)  by  noting  the  brilliancy  of 
the  light  evolved  in  each  case  (three  students  should  work 
together,  each  using  a  different  acid).    Note  the  results. 

b.  From  the  standpoint  of  the  ionization  theory,  upon  what 
does  the  relative  strength  of  two  acids  in  equimolecular  con- 
centration depend  ?  How  will  this  affect  the  rate  at  which 
the  two  acids  will  act  upon  a  third  substance  ?  Will  the  two 
acids,  when  treated  with  a  metal,  evolve  hydrogen  at  the  same 
rate  ?  Will  the  volume  of  hydrogen  ultimately  evolved  be 
the  same  ?  Test  your  statements  by  determining  the  rate  at 
which  each  of  the  three  acids  used  in  a  will  give  up  hydrogen 
when  brought  in  contact  with  zinc.  To  do  this  proceed 
as  follows : 

[54] 


Arrange  a  hydrogen  generator  A  (Fig.  37)  and  fit  it 
with  a  delivery  tube  B.  Fill  a  test  tube  C  with  water  and 
invert  it  in  a  beaker  of  water,  placing  a  slender  rubber  band 
around  it  about  5  cm.  from  the  closed  end.  Weigh  accurately 
2  g.  of  granulated  zinc  and  place  it  in  A.  Through  the  funnel 
tube  add  10  cc.  of  twice  normal  hydrochloric  acid,  to  which 
2  drops  of  cupric  sulfate  has  been  added  (why  ?),  and  bring 
the  end  of  the  delivery  tube 
B  under  C  so  as  to  collect 
the  evolved  hydrogen.  At 
once  note  the  exact  time 
to  the  second.  Keep  the 
rubber  band  at  the  level 
of  the  water  in  the  beaker, 
and  when  the  hydrogen  has 
displaced  the  water  in  C 
to  this  level,  again  read 
the  time  to  the  second.  Re- 
peat the  entire  experiment, 
substituting  twice  normal 
sulfuric  acid  for  the  hy- 
drochloric acid.  Again  re-  FlG  37 
peat,  using  twice  normal 

acetic  acid.  Compare  the  time  required  in  the  three  ex- 
periments. What  do  you  conclude  as  to  the  relative  strengths 
of  the  three  acids  ? 

c.  Repeat  the  three  experiments  in  &,  using  pieces  of  mag- 
nesite  (MgCO3)  of  equal  size  instead  of  the  zinc.  Gaseous 
carbon  dioxide  (CO2)  is  evolved  in  this  reaction.  What  is 
now  the  order  of  the  three  acids  in  strength  ? 

75.  Bases  and  salts  from  the  standpoint  of  the  ionization 
theory,  a.  Bases.  Test  the  conductivity  of  dilute  solutions 
of  the  following  typical  bases:  sodium  hydroxide,  calcium 
hydroxide,  potassium  hydroxide.  Discuss  the  results. 

b.  Salts.  Test  the  conductivity  of  dilute  solutions  of  the 
following  typical  salts:  sodium  sulfate,  ammonium  chloride, 
sodium  acetate.  Discuss  the  results. 

[55] 


76.  Color  of  ions  and  of  molecules.   Prepare  a  small  amount 
of   dilute    solutions    of    each    of    the    following:    potassium 
bromide,  potassium    sulfate,  potassium    chloride.     From  the 
standpoint  of  the  theory  of  ionization,  what  ions  are  present 
in  each  of  the  solutions  ?    What  is  the  color  of  each  of  these 
ions  ?   Place  in  a  test  tube  about  one-half  gram  of  each  of  the 
following  substances:  cupric  bromide,  cupric  sulfate,  cupric 
chloride.    Note  the  color  of  each.    Now  dissolve  each  in  the 
smallest  possible  amount  of  water.    Note  the  color  of  the  solu- 
tions.   Now  gradually  dilute  each  until  the  tubes  are  nearly 
filled  with  water.    Mix  each  thoroughly  and  note  the  color  of 
the  solutions.    Give  an  explanation  of  the  results. 

77.  Electrochemical  series,   a.  Dissolve  about  1  g.  of  copper 
sulfate  in  a  test  tube  two  thirds  full  of  water  and  dip  an 
iron  nail  into  the  solution,  first  rubbing  the  nail  bright  with 
sandpaper  (R).    Reverse  the  experiment  by  dipping  a  bright 
piece  of  copper  foil  into  a  solution  of  ferrous  sulfate.    Is  the 
reaction  reversible  ?    Why  ? 

b.  Suspend  a  piece  of  mossy  zinc  or  tin  by  a  thread  in  a 
solution  of  lead  acetate,  and  let  it  stay  undisturbed,  noting 
the  appearance  of  the  metal  from  time  to  time.    Would  you 
expect  a  piece  of  silver  or  bismuth  to  give  similar  results  ? 

c.  Polish  a  piece  of  aluminium  foil  and  dip  it  for  a  time 
in  pure  water.    Is  hydrogen  evolved  ?    Would  you  expect  it 
to  be  ?    Now  dip  the  foil  into  a  dilute  solution  of  mercuric 
chloride  and  watch  for  evidences  of  evolution  of  hydrogen. 
Is  mercury  deposited  on  the  foil  ?    Would  it  evolve  hydro- 
gen ?    A  deposit  of  mercury  prevents  any  oxide  or  hydroxide 
from  sticking  to  the  aluminium.    Does  this  help  in  explaining 
the  evolution  of  hydrogen? 


[56] 


CHAPTER  XIV 


COMPOUNDS  OF  NITROGEN 

78.  Preparation  of  ammonia,  a.  Dissolve  1  g.  of  ammonium 
chloride  in  a  little  water  in  a  test  tube  and  heat  to  boiling. 
Can  you  detect  the  odor  of  ammonia  ?  Add  a  few  drops  of  a 
solution  of  sodium  hydroxide.  Is  the  ammonia  liberated  (R)  ? 
Moisten  a  small  strip  of  red  litmus  paper  and  hold  it  at  the 
mouth  of  the  tube.  Note  the  result.  Hold  the  end  of  a  glass 
rod  moistened  with  concentrated  hydrochloric  acid  in  the 
mouth  of  the  tube.  What  is  formed  (R)  ? 

b.  Usual  laboratory  method.  This  differs  from  the  method 
used  in  a  only  in  the  fact  that  the  less  expensive  calcium 
hydroxide  (slaked  lime)  is  sub- 
stituted for  sodium  hydrox- 
ide. Fig.  38  shows  the  form  of 
apparatus  used.  A  represents 
a  250-cc.  flask ;  B  and  C  are 
250-cc.  wide-mouthed  bottles 
partly  filled  with  water.  Notice 
that  the  glass  tubes  extending 
into  the  bottles  do  not  quite 
touch  the  surface  of  the  water. 
Why?  Two  dry  bottles  will 
also  be  needed  for  collecting 
some  of  the  gas. 

Put  into  the  flask  A  an  inti- 
mate mixture  of  about  30  g.  of  FlG  38 
finely  powdered    slaked   lime     / 

and  15  g.  of  ammonium  chloride ;  place  it  on  a  sand  bath 
and  heat  gently  with  a  small  flame.    A  wire  gauze  may  be 

[57] 


c: 


substituted  for  the  sand  bath,  provided  the  flask  is  clamped 
so  that  the  bottom  of  it  does  not  quite  touch  the  gauze.  In 
this  case,  however,  the  burner  should  be  held  in  the  hand 
and  moved  about  so  as  to  apply  the  heat  uniformly,  as 
otherwise  the  flask  may  be  broken.  As  soon  as  the  gas  is 
evolved  freely,  disconnect  the  tube  E  at  Z>,  bring  it  to  an 
upright  position,  as  shown  in  the  dotted  lines,  and  collect 
two  bottles  of  the  gas  by  bringing  them  successively  over 
the  exit  tube.  To  tell  when  they  are  filled,  test  for  the  pres- 
ence of  ammonia  at  the  mouths  of  the  bottles  with  a  piece 
of  red  litmus  paper  moistened  with  water.  As  soon  as  the 
bottles  are  filled,  cover  them  tightly  and  set  them  aside, 
mouth  downward ;  then  quickly  connect  the  tube  E  with  the 
bottle  B  again  and  continue  to  heat  the  mixture  gently  as 
long  as  any  gas  is  generated.  What  is  the  source  of  each 
of  the  materials  used  in  the  preparation  of  the  gas  ?  Write 
the  equations  for  all  of  the  reactions  involved. 

79.  Properties  of  ammonia,    a.  Note  the  color  and  odor  of 
the  gas.    Is  it  heavier  or  lighter  than  air  ? 

b.  Test  a  bottle  of  the  gas  with  a  burning  splint.   Describe 
the  results. 

c.  Fill  a  large  beaker  with  water  and  color  it  with  a  few 
drops  of  red  litmus  solution.    Uncover  the  remaining  bottle 
of  the  gas  and  at  once  bring  the  mouth  of  it  under  the  sur- 
face of  the  water  in  the  beaker.    What  do  the  results  prove  ? 

d.  Note  the  odor  of  the  liquid  in  the  bottle  B.    Try  its 
effect  on  blue  and  on  red  litmus  paper.   How  does  it  compare 
with  the  so-called  "  aqua  ammonia  "  of  the  druggist  ?    Does 
the  gas  combine  with  the  water,  or  is  it  simply  dissolved  in 
it  ?   Give  reasons  for  your  answer.   Now  neutralize  the  liquid 
with  hydrochloric  acid  (R)  and  evaporate  just  to  dryness. 
What  is  the  residue  ?    Test  it  with  a  few  drops  of  sodium 
hydroxide  solution.    How  does  it  compare  with  one  of  the 
compounds  used  in  the  preparation  of  the  ammonia?    Dis- 
tinguish between  the  terms  ammonia  and  ammonium. 

e.  Advanced  students  may  at  this  time  determine  the  com- 
position of  ammonia  according  to  §  114. 

[58] 


80.  Nitrides.  Connect  the  delivery  tube  E  (Fig.  38)  with 
the  hard-glass  tube  C  (Fig.  26).  In  the  middle  of  the  tube 
place  a  shallow  layer  of  magnesium  powder  about  2  cm.  in 
length.  Conduct  ammonia  through  the  tube,  and  when  the 
air  has  been  displaced,  heat  the  tube  under  the  magnesium 
strongly  until  the  magnesium  is  red  hot.  Then  increase 
the  current  of  ammonia.  What  change  takes  place  in  the 


FIG.  39 

magnesium  (R)  ?  When  it  is  cool  examine  the  powder.  Place 
a  little  of  it  on  a  watch  glass  and  moisten  it  with  water.  Can 
you  detect  the  evolution  of  ammonia  (R)  ? 

81.  Weight  of  i  liter  of  ammonia.    Arrange  an  apparatus 
as  shown  in  Fig.  39.    A  is  a  250-cc.  flask  containing  about 
100  cc.   of  concentrated  ammonia  water.    B  is  a  glass  tube 
entirely  filled  with  pieces  of  calcium  oxide  (unslaked  lime) 
about  the  size  of  peas.    C  is  the  flask  used  in  determining  the 
weight  of  1  liter  of  carbon  dioxide  (Fig.  30).    Heat  the  am- 
monia water  gently.    Ammonia  is  evolved,  is  dried  in  passing 
over  the  lime  in  B,  and  collects  in  C.    After  C  is  filled  with 
ammonia,  proceed  as  in  determining  the  weight  of  1  liter  of 
carbon  dioxide.    Give  the  results  obtained. 

82.  Preparation  of  nitric  acid.    a.  Arrange  an  apparatus  as 
shown  in  Fig.  40.    The  retort  A  should  have  a  capacity  of 
from  100  to  150  cc.    Put  into  the  retort  about  12  g.  of  sodium 

[59] 


nitrate  and  10  cc.  of  concentrated  sulfuric  acid,  pouring  the 
latter  through  a  funnel  placed  in  the  tubulus  of  the  retort. 
Heat  the  mixture  gently  with  a  small  flame.  Nitric  acid  is 
generated  (R),  distills  over,  and  is  condensed  in  the  test 
tube  B,  which  should  be  kept  cool  with  water.  How  many 
grams  of  nitric  acid  can  be  prepared  from  the  amount  of 
materials  used  ? 

b.  What  is  the  color  of  the  acid  you  have  prepared  ?  What 
is  the  color  of  the  pure  acid  ?  How  do  you  account  for  the 
difference  (R)  ?  Try 
heating  a  little  of  the 
colored  product  to  gen- 
tle boiling.  How  does 
this  affect  the  color  ? 
Why? 

83.  Action   of  nitric 
acid  on  metals,  a.  Place 


one  or  two  pieces  of 
mossy  tin  in  a  test  tube 
and  add  2  or  3  cc.  of 
the  nitric  acid  prepared 
in  §  82.  Note  the  color  of  the  evolved  gas  during  the  re- 
action. Does  the  gas  in  the  test  tube  remain  colored,  or  does 
the  brownish  color  soon  disappear  ?  What  is  the  chief  gas 
evolved  under  these  conditions  ?  When  the  action  is  over, 
dilute  to  about  three  times  the  volume.  The  white  precipitate 
is  a  hydrate  of  the  oxide  SnO2,  and  for  the  present  purpose 
it  may  be  regarded  as  the  oxide  itself.  Filter  off  the  solid 
material,  carefully  collecting  the  nitrate.  If  a  nitrate  of  tin 
has  been  formed,  it  should  be  in  the  nitrate,  for  it  is  soluble 
in  water.  Add  an  excess  of  ammonium  hydroxide  to  the 
nitrate.  Is  a  precipitate  formed  ?  Any  tin  present  in  the  ni- 
trate will  be  precipitated  as  the  hydroxide.  Verify  this  state- 
ment by  adding  a  little  ammonium  hydroxide  to  a  solution  of 
tin  chloride.  Is  tin  nitrate  formed  by  the  action  of  concen- 
trated nitric  acid  on  the  metal?  Formulate  an  equation  for 
the  chief  reaction,  assuming  that  the  white  solid  is  SnO2. 

[60] 


FIG.  40 


b.  Dilute  a  little  of  the  concentrated  nitric  acid  with  three 
times  its  volume  of  water,  and  pour  the  dilute  acid  upon  a 
few  pieces  of  copper  turnings  in  a  test  tube.    What  is  the 
color  in  the  test  tube  as  the  action  begins  ?  What  change  in 
color  do  you  notice  ?  Hold  a  piece  of  white  paper  behind  the 
test  tube  for  a  background,  and  compare  the  color  of  the  gas 
at  the  mouth  of  the  tube  with  that  within  the  tube.    How  do 
you  account  for  all  the  facts  observed  ?  What  is  the  ch'ief  gas 

•*  o 

evolved  in  this  reaction  (R)  ? 

c.  Place  one  or  two  pieces  of  tin  in  a  test  tube,  add  5  cc.  of 
water,  and  then  from  time  to  time  add  a  drop  of  concentrated 
acid.    Keep  shaking  the  solution,  and  continue  the  operation 
until  you  have  added  at  least  1  cc.  of  nitric  acid.    Filter  the 
solution  and  add   ammonium  hydroxide  in  excess.    Does  a 
precipitate  form  ?  What  does  this  indicate  ?  Has  any  notice- 
able quantity  of  red  gas  been  evolved  ?   From  its  position 
in  the  electromotive  series  how  would  you  expect  tin  to  act 
upon  dilute  nitric  acid  (R)  ? 

d.  Place  a  few  pieces  of  zinc  in  a  test  tube  and  add  dilute 
sulfuric  acid  until  a  gas  is  freely  evolved.   What  is  it?   From 
time  to  time  add  a  drop  of  concentrated  nitric  acid.    What 
action  would  you  expect  to  take  place  between  nascent  hy- 
drogen and  nitric  acid?    If  ammonia  is  formed,  what  would 
become  of  it  (remember  what  is  present  in  the  solution)  ? 
How  could  you  test  for  its  presence  (R)  ?    Can  you  detect  it 
in  the  present  experiment? 

e.  Give  a  summary  of  how  nitric  acid,  both  concentrated  and 
dilute,  may  be  expected  to  act  upon  metals  above  hydrogen  in 
the  electromotive  series,  and  also  upon  those  below  hydrogen. 

84.  Salts  of  nitric  acid :  nitrates,  a.  Place  a  small  piece 
of  copper  in  an  evaporating-dish  (hood),  add  a  few  drops  of 
the  ordinary  nitric  acid,  and  heat  gently.  As  soon  as  the 
copper  has  dissolved,  carefully  evaporate  the  solution.  The 
flame  must  not  touch  the  dish,  and  should  be  withdrawn 
while  two  or  three  drops  of  the  liquid  still  remain  in  the 
dish.  What  is  the  composition  of  the  residue  (R)  ?  Apply 
a  gentle  heat  to  the  residue  left  in  the  evaporating-dish. 

[61] 


Account  for  the  change  in  color  (R).  Repeat  the  experi- 
ment, using  a  piece  of  zinc  or  lead  instead  of  copper.  Are 
all  nitrates  decomposed  by  heat? 

b.  Test  the  solubility  in  water  of  sodium  nitrate,  potassium 
nitrate,  barium  nitrate,  lead  nitrate,  and  copper  nitrate,  using 
a  small  crystal  in  each  case.    Describe  the  results.    What 
nitrates  are  insoluble  ?    Try  a  crystal  of  bismuth  nitrate  or 
ferric  nitrate. 

c.  Dissolve  a  crystal  of  sodium  nitrate  in   2  or  3  cc.  of 
water  in  a  test  tube,  add  an  equal  volume  of  sulfuric  acid, 
and  cool.    Now  incline  the  tube  and  gently  pour  a  concen- 
trated solution  of  ferrous  sulfate  down  the  side  of  the  tube, 
so  that  it  floats  on  the  heavier  liquid.    A  brown  ring  soon 
forms  where  the  liquids  meet.    Repeat  the  experiment,  using 
potassium  nitrate.    This  is  a  good    test  for  nitrates.    The 
brown  ring  is  due  to  the  formation  of  a  compound  of  fer- 
rous sulfate  and  nitric  oxide,  the  latter  being  formed  by  the 
reduction  of  the  nitric  acid  by  the  ferrous  sulfate. 

85.  Combining  weights  of  metals  from  oxidation  to  oxides. 
The  following  exercise  will  serve  as  an  illustration  of  a  gen- 
eral method.  Weigh  accurately  a  small  evaporating-dish  or  a 
large  crucible,  place  in  it  about  1  g.  of  pure  copper,  and  again 
weigh.  Cover  the  dish  with  a  watch  glass  and  add  dilute 
nitric  acid,  a  little  at  a  time,  until  the  metal  is  all  dissolved. 
Rinse  the  watch  glass  carefully  into  the  dish,  and  place  the 
latter  on  the  rim  of  a  somewhat  smaller  beaker  two  thirds 
full  of  water.  By  boiling  the  water  in  the  beaker,  evaporate 
the  solution  of  copper  nitrate  until  it  is  nearly  dry.  Place 
the  dish  on  a  pipestem  triangle  and  very  carefully  heat  it 
with  a  small  flame,  holding  the  burner  in  the  hand  and  using 
every  precaution  to  prevent  loss  by  spattering.  As  the  nitrate 
becomes  dry,  gradually  increase  the  heat  until  red  fumes  cease 
to  be  given  off,  finally  using  the  full  heat  of  the  burner.  Allow 
the  dish  to  cool  and  then  weigh  it.  What  is  the  product? 
From  the  increase  in  weight  calculate  the  ratio  by  weight  in 
which  copper  and  oxygen  combine.  The  experiment  may  be 
varied  by  the  use  of  iron,  tin,  or  zinc  in  place  of  copper. 

[62] 


86.  Preparation  and  properties  of  nitrous  acid.    In  a  hemi- 
spherical iron  dish  heat  10  g.  of  potassium  nitrate  until  it 
melts  and  just  begins  to  evolve  bubbles ;  then  add  25  g.  of 
lead.    Continue  the  heating  for  about  twenty  minutes,  stirring 
the  mixture  with  an  iron  wire  or  file.    Note  the  change  in 
color.    How  do  you  account  for  it  (R)  ?    When  the  product 
becomes  cool  add  25  cc.  of  water  and  heat  until  the  mass  is 
disintegrated.    Filter  off  the  residue.    What  is  its  composi- 
tion ?    What  compound  is  present  in  the  filtrate  ?    Add  to 
the  filtrate  a  few  drops  of  sulfuric  acid.     Account  for  the 
result.    Write  the  equation  for  the  reaction  which  takes  place 
between  sulfuric  acid  and  potassium  nitrate ;  between  sulfuric 
acid  and  potassium  nitrite. 

87.  Preparation  and  properties  of  some  of  the  oxides  of  nitro- 
gen,   a.  Nitrous  oxide.    Put  6  or  8  g.  of  ammonium  nitrate  in 
the  hard-glass  test  tube  used  in  the  preparation  of  oxygen. 
Attach  a  delivery  tube  and  heat  gently,  applying  no  more  heat 
than  is  necessary  to  cause  a  slow  evolution  of  gas.    As  soon 
as  the  gas  is  regularly  evolved  collect  two  or  three  bottles  of 
it  over  water.    Notice  the  deposit  of  water  on  the  sides  of  the 
test  tube.    What  is  the  source  of  it  (R)  ?    Note  the  color, 
odor,  and  taste  of  the  gas.    Test  it  with  a  glowing  splint. 
Account  for  the  result.    How  can  you  distinguish  it  from 
oxygen  ?    What  is  the  common  name  of  the  gas,  and  for 
what  is  it  used  ?    Contrast  the  action  of  heat  on  ammonium 
nitrate  and  copper  nitrate  (R). 

b.  Nitric  oxide  and  nitrogen  dioxide.  Put  a  few  pieces  of 
copper  in  your  hydrogen  generator  (hood),  just  cover  them 
with  water,  and  add  2  or  3  cc.  of  nitric  acid.  Collect  over 
water  three  bottles  of  the  evolved  gas,  adding  more  nitric 
acid  if  necessary,  and  leaving  the  last  one  half  filled  in  the 
pneumatic  trough.  Compare  the  color  of  the  gas  in  the  gen- 
erator with  that  collected  in  the  bottles  and  account  for  any 
difference.  Write  the  equations  for  all  the  reactions  involved. 
Uncover  one  of  the  bottles  containing  gas  and  account  for 
the  result.  Test  the  gas  in  the  second  bottle  with  a  burning 
splint.  Which  is  the  more  stable,  nitrous  oxide  or  nitric 

[63] 


oxide  ?  Give  reasons  for  your  answer.  To  the  third  bottle 
standing  over  water  add  air  in  small  portions  at  a  time, 
transferring  the  air  to  the  bottle  with  a  test  tube.  After  the 
addition  of  each  portion  of  air,  allow  the  red  fumes  to  dis- 
solve before  adding  another  portion.  Does  the  volume  dimin- 
ish indefinitely  ?  Why  ?  If  pure  oxygen  had  been  added 
instead  of  air,  would  all  the  gas  in  the  bottle  vanish?  Can 
you  suggest  a  modification  of  the  experiment  that  could  be 
used  to  determine  the  percentage  of  oxygen  in  the  air? 


[64] 


CHAPTER  XV 

EQUILIBRIUM 

88.  Velocity  of  reactions,    a.  To  a  dilute  solution  of  sodium 
chloride  add  1  or  2  cc.  of  a  solution  of  silver  nitrate.    How 
rapidly  does  the  reaction  take  place  ?     How  does  it  compare 
in  rapidity  with  the  reaction  occasioned  by  the  addition  of  a 
little  barium  chloride  to  dilute  sulfuric  acid  ?   In  general,  re- 
actions between  freely  ionized  electrolytes  are  too  rapid  for 
measurement.     Slow  reactions  occur  when  the  reagents  are 
very  little  ionized  or  when  changes  occur  other  than  double 
decomposition  between  ions. 

b.  Dissolve  a  crystal  of  sodium  phosphate  no  larger  than 
a  grain  of  wheat  in  10  cc.  of  water  and  add  5  cc.  of  ammonium 
molybdate  solution  (side  shelf).  Does  a  precipitate  at  once 
form  ?  Do  you  notice  any  change  of  color  ?  Does  this  re- 
main constant  or  increase  ?  Set  the  tube  aside  and  look  at 
it  from  time  to  time  during  the  laboratory  period.  Is  the 
formation  of  a  precipitate  gradual  or  sudden  ?  (The  yellow 
solid  has  a  very  complex  formula.) 

89.  Mass  action,    a.  Make  a  concentrated  solution  of  com- 
mon salt  by  shaking  20  g.  of  salt  with  the  least  water  that 
will  serve  to  dissolve  it.    Add  5  cc.  of  concentrated  sulfuric 
acid  in  small  portions  at  a  time,  shaking  gently  after  each 
addition  (hood).    The  gas  evolved  has  the  formula  HC1,  the 
equation  being 

2  NaCl  +  H2SO4  =  Na2SO4  +  2  HC1 

When  all  of  the  acid  has  been  added,  warm  gently  and  set 
the  solution  aside  to  crystallize.  (If  no  crystals  appear  by 
the  time  the  solution  has  cooled,  add  a  minute  crystal  of 

[65] 


sodium  sulfate  (Na2SO4).)  How  do  the  crystals  differ  in  ap- 
pearance from  those  of  common  salt?  Dry  a  few  of  them 
and  expose  them  to  the  air  (R).  Warm  a  few  in  a  test  tube. 
Do  they  contain  water  of  hydration  ?  Does  salt  ?  Drain  the 
mother  liquor  from  the  crystals  and  pour  over  them  about 
5  cc.  of  concentrated  hydrochloric  acid.  Wnat  members  of 
the  above  equation  have  you  brought  together  ?  Do  you 
note  any  change  in  the  crystals  ?  Filter  some  of  them  and 
dry  them  on  filter  paper.  How  can  you  prove  that  they  are 
common  salt  ?  Complete  the  equation.  How  does  it  com- 
pare with  the  one  written  above  ?  Why  has  the  reaction 
been  reversed  ? 

b.  Place  a  few  crystals  of  bismuth  chloride  (BiCl3)  in  a 
beaker  and  add  about  1  cc.  of  water.  What  change  do  you 
notice  ?  The  equation  is 

BiCl3  +  H20  =  BiOCl  +  2  HC1 

Under  what  conditions  would  you  expect  the  reaction  to  be 
reversed  ?  Add  concentrated  hydrochloric  acid,  a  drop  at  a 
time,  with  constant  stirring.  What  is  observed  (R)?  When 
the  solution  is  complete,  add  water,  a  little  at  a  time,  until 
the  precipitate  again  appears.  Can  you  again  reverse  the 
reaction  ?  How  many  times  can  you  repeat  the  process  ? 

90.  Equilibrium  in  solution,  a.  Dissolve  about  1  g.  of  cupric 
bromide  (CuBr2)  in  the  least  possible  volume  of  water.  What 
is  the  color  of  the  solution  ?  Add  water,  a  small  quantity  at 
a  time,  until  the  solution  becomes  a  clear  green.  Pour  off 
half  the  solution  and  dilute  it  further.  What  is  the  final 
color  ?  How  can  you  account  for  the  intermediate  green  ? 
To  the  remaining  half  of  the  green  solution  add  crystals  of 
potassium  bromide  or  sodium  bromide.  Yfhat  change  in  color 
is  noticed  ?  How  do  you  explain  this  ? 

b.  Obtain  about  5  cc.  of  a  solution  of  ammonium  sulfo- 
cyanate  (NH4CNS)  and  of  ferric  chloride  (FeCl8)  from  the 
side  shelf.  Pour  1  cc.  of  each  of  these  solutions  into  50  cc. 
of  water.  An  equilibrium  results  as  follows: 

3  NH4CNS  +  FeCl3  +=±  Fe(CNS)3  +  3  NH4C1 
[66] 


The  compound  Fe(CNS)3  is  deep  red,  and  even  in  very  dilute 
solution  has  a  noticeable  pink  color.  Half  fill  four  test  tubes 
with  the  solution,  and  reserve  one  tube  for  comparison.  To 
the  second  tube  add  1  or  2  cc.  of  the  solution  of  ammonium 
sulfocyanate.  What  change  in  color  results  ?  Note  the  equa- 
tion and  explain.  To  the  third  tube  add  1  or  2  cc.  of  the 
solution  of  ferric  chloride.  Explain  the  change.  To  the  fourth 
tube  add  2  or  3  g.  of  solid  ammonium  chloride.  Explain  the 
change. 


[67] 


CHAPTER  XVI 

SULFUR  AND  ITS  COMPOUNDS 

91.  Properties  of  sulfur,  a.  Examine  the  physical  proper- 
ties of  a  piece  of  brimstone.  Pour  2  or  3  cc.  of  carbon  disul- 
fide  over  a  little  powdered  brimstone  in  a  test  tube  (keep 
away  from  flame).  Cover  the  mouth  of  the  tube  with  the 
thumb  and  shake  the  contents  gently  until  the  sulfur  is  dis- 
solved. Pour  the  clear  solution  into  an  evaporating-dish, 
cover  it  loosely  with  a  filter  paper,  and  set  it  aside  in  the 
hood.  The  carbon  disulfide  soon  evaporates,  the  sulfur  being 
deposited  in  crystals.  Examine  these  with  a  magnify  ing-glass. 

b.  Half  fill  a  test  tube  with  powdered  brimstone  and  apply 
heat  enough  to  just  melt  it.   Note  the  properties  of  the  liquid. 
Pour  a  little  of  the  liquid  into  a  beaker  of  water,  dry  the 
product  on  filter  paper,  and  test  its  solubility  in  carbon  disul- 
fide.   Now  apply  a  stronger  heat  and  observe  that  the  liquid 
becomes  darker  and  at  a  certain  temperature  (200°- 250°)  is 
so  thick  that  the  tube  may  be  inverted  without  spilling  it. 
Finally,  increase  the  heat  until  the  sulfur  boils  (445°),  and 
pour  the  boiling  liquid  into  a  beaker  of  cold  water.    Examine 
the  product.     What  name  is  given  to  this  form  of  sulfur? 
Dry  a  small  piece  on  filter  paper  and  test  its  solubility  in 
carbon  disulfide.   Expose  it  to  the  air  for  a  day  or  two.  Have 
its  properties  remained  unchanged  ? 

c.  Fill  a  porcelain  crucible  with  powdered  brimstone  and 
apply  a  very  gentle  heat  until  the   sulfur  is  just  melted. 
Withdraw  the   flame   and  watch  the   liquid  carefully  as  it 
cools.    Crystals  soon  begin  to  form  on  the  surface,  rapidly 
growing  from  the  circumference  toward  the  center.    Before 
they  reach  the  center  quickly  pour  the  remaining  liquid  into 

[68] 


a  dish  and  examine  the  crystals  adhering  to  the  sides  of  the 
crucible.  Compare  them  with  those  formed  in  a.  In  how 
many  forms  have  you  obtained  the  sulfur?  Which  is  the 
stable  form  ? 

d.  Burn  a  small  piece  of  sulfur.  What  is  the  product  of  the 
combustion  ?  Note  the  odor  of  it.  Heat  to  boiling  a  little 
sulfur  in  the  test  tube  used  in  b  and  drop  a  small  strip  of 
bright  copper  foil  into  the  boiling  liquid.  Is  there  any  visible 
evidence  of  a  chemical  reaction  ?  What  is  formed  ?  What  is 
the  product  formed  in  §  11,  b? 

92.  Transition  point.  The  transition  of  monoclinic  sulfur 
into  rhombic  is  too  slow  to  be  observed  readily.  The  following 
is  a  more  striking  experiment :  Obtain  a  piece  of  thin-walled 
glass  tubing  about  15  cm.  long  and  5  mm.  in  diameter.  Seal 
one  end  in  the  Bunsen  flame,  blowing  the  seal  into  rounded 
form  like  a  test  tube.  Obtain  about  half  a  gram  of  cupro- 
mercuric  iodide  (side  shelf)  and  put  it  into  this  test  tube. 
Half  fill  a  beaker  with  water,  place  it  on  a  wire  gauze  on 
a  ring  stand,  and  suspend  a  thermometer  in  such  a  way  that 
the  bulb  dips  into  the  water.  Heat  the  water  until  the  tem- 
perature reaches  about  60°,  then  turn  the  flame  down  so  that 
the  rise  in  temperature  is  very  slow.  Use  the  test  tube  con- 
taining the  scarlet  powder  as  a  stirring-rod  and  watch  for  a 
change  in  color.  What  change  occurs  ?  At  what  tempera- 
ture ?  Withdraw  the  flame,  allowing  the  temperature  to 
fall  slowly.  At  what  point  does  the  reverse  change  occur  ? 
Repeat  the  experiment  to  ascertain  within  what  limits  of 
temperature  the  change  occurs.  Return  the  powder  to  the 
stock  bottle.  Cupromercuric  iodide  is  sometimes  used  as  a 
paint  for  bearings  in  machinery,  since  its  change  in  color 
indicates  that  the  bearing  is  becoming  heated. 

NOTE.  If  desired,  the  student  may  himself  prepare  the  cupromercuric 
iodide  as  follows :  Dissolve  a  few  crystals  of  mercuric  chloride  in  2  or 
3  cc.  of  hot  water  in  one  test  tube,  a  few  crystals  of  potassium  iodide 
in  another,  about  the  same  weight  of  copper  sulfate  in  a  third,  and' 
a  similar  quantity  of  sodium  sulfite  in  a  fourth.  Add  the  potassium 
iodide,  drop  by  drop,  to  the  mercuric  chloride  until  the  bright  scarlet 

[69] 


precipitate  just  redissolves.  To  the  solution  of  sodium  sulfite  add  a 
little  moderately  dilute  sulfuric  acid,  and  pour  this  solution  into  the 
one  just  prepared;  then  quickly  add  the  copper  sulfate.  Collect  the  red 
cupromercuric  iodide  on  a  filter  paper,  wash  it  once  or  twice  with  cold 
water,  and  dry  the  powder  on  sheets  of  filter  paper. 

93.  Hydrogen  sulfide.  a.  (Hood.)  Attach  a  delivery  tube 
to  the  hydrogen  generator,  as  represented  in  Fig.  41.  Put 
into  the  generator  a  few  pieces  of  iron  sulfide  (FeS)  and 
insert  the  stopper.  Before  generating  the  gas  have  at  hand 
a  dry  bottle  in  which  to  collect  some  of  it,  a  short  piece  of 
glass  tubing  of  size  to  fit  the  rubber  connection  A  and  drawn 
to  a  jet  at  one  end,  also  test 
tubes  about  one  fifth  full  of 
solutions  of  each  of  the  fol- 
lowing compounds :  copper 
sulfate,  zinc  sulfate,  cadmium 
sulfate,  magnesium  sulfate, 
sodium  chloride.  Now  pour 
a  few  drops  of  a  dilute  solu- 
tion of  sulfuric  acid  (1  part 
acid  to  7  parts  of  water  by 
volume)  through  the  funnel 
tube  of  the  generator,  adding 
more  from  time  to  time,  if 

necessary,  to  produce  a  gentle  evolution  of  the  gas  (R). 
Be  careful  not  to  inhale  the  gas,  as  it  is  poisonous. 

b.  Expose  a  piece  of  moist  litmus  paper  to  the  gas.    What 
is  its  reaction  ?    Dip  a  piece  of  filter  paper  into  a  solution  of 
lead  acetate  (side  shelf)  and  expose  it  to  the  gas  (R). 

c.  Collect  a  bottle  of  the  gas  by  downward  displacement, 
allowing  the  gas  to  flow  until  it  is  ignited  by  a  flame  held  at 
the  mouth  of  the  bottle  (R).    Account  for  the  deposit  on  the 
sides  of  the  bottle.     Connect  the  glass  jet  with  the  delivery 
tube  at  A  and  ignite  the  gas  at  the  tip.    Note  the  appear- 
ance of  the  resulting  flame.    What  are  the  products  of  the 
combustion  ?    Hold  a  cold  porcelain  dish  in  the  flame.    What 
is  deposited  on  the  dish  ?    Why  ? 

[70] 


d.  Cause  the  gas  to  bubble  for  a  few  seconds  through  each  of 
the  solutions  in  the  test  tubes  (Fig.  41),  noting  the  color  of  the 
precipitates  obtained.     What  is  the  composition  of  each  (R)  ? 
In  which  cases  does  no  precipitate  form  ?  Why  is  this  ? 

e.  Pass  a  little  of  the  gas  into  some  water  in  a  beaker. 
How  does  the  solution  compare  with  natural  sulfur  waters  ? 
Note  the  odor  of  the  solution.    Pour  a  little  of  it  into  a  test 
tube  and  boil  for  a  short  time.    Does  the  odor  disappear? 
Can  the  gas  be  expelled  from 

the  water  by  boiling?  Drop 
a  silver  coin  into  the  solution, 
and  account  for  the  results. 

/.  Pass  the  gas  into  a  solu- 
tion of  zinc  sulfate  as  long  as  a 
precipitate  continues  to  form. 
If  the  action  is  a  completed 
one,  will  any  zinc  sulfate  re- 
main in  the  solution  ?  Filter 
from  the  zinc  sulfide  and  add 
a  little  ammonium  hydroxide 
to  the  filtrate.  How  do  you 
account  for  the  additional 
precipitate  ? 

94.  Sulfur  dioxide.  Place  about  10  g.  of  copper  turnings 
or  small  pieces  of  sheet  copper  in  a  generator  arranged  as  in 
Fig.  42  (hood).  Add  about  25  cc.  of  concentrated  sulfuric  acid 
and  apply  a  gentle  heat.  As  soon  as  the  action  begins  (R), 
lower  the  flame,  regulating  it  so  as  to  obtain  a  uniform  evo- 
lution of  the  gas.  Collect  two  bottles  of  the  gas  by  downward 
displacement ;  then  cause  it  to  bubble  through  25  cc.  of  water 
as  long  as  any  is  dissolved.  Do  not  allow  the  delivery  tube  to 
remain  in  the  solution  after  the  gas  ceases  to  flow,  or  the 
solution  will  run  back  into  the  flask  (why  ?).  By  what  other 
method  has  this  gas  been  prepared  in  a  previous  exercise  ? 
Invert  one  of  the  bottles  of  the  gas  so  that  the  mouth  of  it  is 
under  water,  and  examine  the  contents  after  several  minutes. 
In  the  second  bottle  place  some  moist  paper  containing  writing 

[71] 


FIG. 42 


in  ink,  some  paper  with  print  upon  it,  or  some  moistened 
flowers  or  colored  cloth.  Are  the  colors  affected?  Contrast 
the  action  of  hot  concentrated  sulfuric  acid  upon  copper  with 
that  of  dilute  sulfuric  acid  on  zinc  or  iron  (R). 

95.  Sulfurous  acid.    a.  Test  the  saturated  aqueous  solution 
of  sulfur  dioxide  with  blue  litmus.    Is  the  gas  combined  with 
the  water  or  simply  dissolved  in  it?    Give  reason  for  your 
answer.     Boil  a  little  of  the  solution  vigorously  for  a  time. 
Does  the  odor  persist  ?    How  do  you  account  for  this  ?    To 
some  of  the  concentrated  solution  add  a  solution  of  sodium 
hydroxide,    drop    by   drop,    until   it    becomes   neutral  (R), 
and  then  evaporate  just  to  dryness.   What  is  the  residue  ? 
Moisten  it  with  2  or  3  drops  of  sulfuric  acid  and  note  the 
odor  of  the  gas  evolved.    What  is  it?    How  could  you  test 
for  a  sulfite? 

b.  To  5  cc.  of  the  solution  of  sulfur  dioxide  add  a  drop  or 
two   of   barium  chloride.    Does  a  precipitate    form?    Is    it 
soluble  in  hydrochloric  acid?    Repeat  the  experiment,  first 
adding  a  little  concentrated  nitric  acid,  and  boiling  before 
adding  the  barium  chloride.    Do  you  note  any  differences? 
How  do  you  explain  them? 

c.  Place  about  10  cc.  of  water  in  a  bottle  and  saturate  it 
with  sulfur  dioxide  ;  then  pass  hydrogen  sulfide  into  the  solu- 
tion.   What  is  the  precipitate  (R)?    Can  you  filter  it  off? 
What  is  this  form  called? 

96.  Weight  of  i  liter  of  sulfur  dioxide.   Determine  the  weight 
of  1  liter  of  sulfur  dioxide  according  to  the  method  suggested 
for  carbon  dioxide  (§  52),  making  the  necessary  alterations  in 
the  method.    Will  it  be  necessary  to  dry  the  gas  ?    Why  ? 

97.  Sulfuric  acid.    a.   Usual  laboratory  method  (hood).    Ar- 
range an  apparatus  according  to  Fig.  43,  in  which  B  represents 
a  wide-mouthed  bottle  of  about  1  liter  capacity.    The  bottle 
should  be  rinsed  out  with  water  but  not  dried.    By  the  action 
of  concentrated  nitric  acid  on  copper,  generate  nitrogen  diox- 
ide in  C  until  B  is  filled  with  red  fumes.    Now  pass  in  sulfur 
dioxide  from  A  until  the  red  fumes  entering  from  C  are  com- 
pletely decolorized ;  then  replace  A  with  a  flask  containing 

[72] 


a  little  water,  boil  the  water,  and  conduct  the  steam  into  B 
until  5  or  10  cc.  of  liquid  has  been  collected.  Pour  this  into 
a  beaker  and  save  it  for  further  experiments.  Complete  the 
following  equations  : 

Cu  +  HN03  = 
Cu  +  H2SO4  = 


b.  Add  3  drops   of  sulfuric  acid  to  5  cc.   of  water  in  a 
test  tube.    To  this  add  a  few  drops  of  a  solution  of  barium 
chloride.  Note  that 

a  precipitate  forms 

(R).  Now  add  3  or 

4  drops  of  hydro- 

chloric acid.    Does 

the  precipitate  dis- 

solve ?  The  forma- 

tion   with    barium 

chloride  of  a  pre- 

cipitate    which     is 

insoluble  in  hydro- 

chloric acid  consti- 

tutes   a  good  test  ElG  43 

for    sulfuric    acid. 

Now  apply  this  test  to   the  liquid  taken  from  the  bottle  B 

(Fig.  43).    Note  the  results. 

c.  Recall  the  action  of  dilute  sulfuric  acid  on  zinc  (R). 
Try  the  action  of  the  concentrated  acid  on  zinc,  applying  a 
gentle  heat,  if  necessary,  to  start  the  reaction.    What  gas  is 
now  evolved  ?   Explain.   Write  in  steps  the  equations  for  the 
reactions  which  take  place  when  copper  is  acted  on  by  nitric 
acid  and  hot  concentrated  sulfuric  acid  respectively,  pointing 
out  the  similarity  in  the  action  of  the  two  acids.    Gently  heat 
a  small  bit  of  charcoal  with  2  or  3  cc.  of  concentrated  sulfuric 
acid.    What  gas  is  evolved  (R)  ? 

[73] 


d.  Put  a  drop  of  concentrated  sulfuric  acid  on  a  splint  and 
gently  warm  it  above  a  flame.    Pour  a  few  drops  on  0.5  g.  of 
sugar  in  a  test  tube.    Examine  after  a  few  minutes.    Account 
for  the  results. 

e.  Why  is  sulfuric  acid  used  in  the  preparation  of  other 
acids  ? 

98.  Salts  of  sulfuric  acid :  sulfates.  a.  Obtain  a  crystal  of 
each  of  five  or  six  different  sulfates,  dissolve  them  in  water, 
and  test  each  solution  with  litmus  paper.  Why  do  some  of 
the  sulfates  have  an  acid  reaction  ?  Try  to  dissolve  a  crystal 
of  sulfate  of  antimony,  of  bismuth,  or  of  mercury  in  pure 
water.  What  do  you  observe  ?  Explain. 

b.  Dissolve  in  a  little  water  a  crystal  of  each  of  the  fol- 
lowing sulfates :  sodium  sulfate,  magnesium  sulfate,  copper 
sulfate.    Apply  to  each  the  barium  chloride  test  for  sulfuric 
acid  (R).    Can  sulfates  be  detected  by  this  method?    Dis- 
solve a  crystal  of  sodium  sulfite  in  water  and  apply  the  same 
test.  How  can  you  distinguish  between  sulfates  and  sulfites  ? 
(Sulfites  very  often  contain  more  or  less  of  the  correspond- 
ing sulfates,  owing  to  the   absorption  of   oxygen  from  the 
atmosphere.) 

c.  Is  sulfuric  acid  monobasic  or  dibasic  ?    Explain  these 
terms.    With  due  care  pour  about  10  cc.  of  concentrated  sul- 
furic acid  into  100  cc.  of  water,  and  divide  the  solution  into 
two  equal  parts.  Carefully  neutralize  the  one  part  with  sodium 
hydroxide  and  evaporate  the  solution  to   about  25  cc.    Set 
aside  for  a  day  to  obtain  crystals.    What  are  they  ?    Divide 
the  second  part  into  two  equal  portions.    Neutralize  the  one 
part  as  before ;  then  add  the  second  portion.    What  should 
be  in  solution  ?  Evaporate  as  in  the  other  case  and  set  aside 
to  crystallize.    Compare  the  two  kinds  of  crystals  in  appear- 
ance.   Heat  a  few  of  each  kind  in  a  dry  test  tube.    What 
difference  do  you  note  ? 


[74] 


CHAPTER  XVII 
THE  CHLORINE  FAMILY 

NOTE.  The  experiments  on  chlorine,  hydrogen  chloride,  and  hydro- 
chloric acid  have  been  given  in  Chapter  X.  The  student  should  review 
these  in  connection  with  the  present  chapter. 

99.  Hydrogen  fluoride;  hydrofluoric  acid.    (The  gas  is  very 
corrosive , and  must  not  be  inhaled;  its  solution  must  not  be 
brought  in  contact  ivith  the  skin.')    Warm  gently  over  a  small, 
luminous  flame  a  glass  plate  on  which  some  pieces  of  paraffin 
have  been  placed.    When  the  paraffin  is  melted,  tilt  the  plate 
about  so  as  to  cover  it  with  a  uniform  layer  of  the  wax. 
When   the  plate  becomes   cold,  scratch  your  name  through 
the  wax  with  a  pin  or  other  sharp  point.    Place  in  a  lead  dish 
3  g.  of  powdered  fluorite  and  add  sufficient  sulfuric  acid  to 
make  a  paste  of  it.    Cover  the  dish  tightly  with  the  waxed 
side  of  the  glass  plate  and  set  it  in  the  hood  for  an  hour; 
then  remove  the  paraffin  and  examine  the  glass.    Write  the 
equations  for  all  the  reactions  involved.    In  what  kind  of 
bottles  is  hydrofluoric  acid  stored  ? 

100.  Salts  of  hydrochloric  acid :  chlorides.    Place  in  separate 
test  tubes  a  few  drops  of  a  solution  of  each  of  the  following 
compounds:    sodium    chloride,    potassium    chloride,   calcium 
chloride.    Add  to  each  1  or  2  drops  of  a  solution  of  silver 
nitrate.     Explain    the    results    (R).     Add    a    few    drops   of 
ammonium  hydroxide  to  each  test  tube,  and  note  the  result. 
Now  add  nitric  acid,  drop  by  drop,  to  each  test  tube  until  the 
liquid  is  acid,  and  note  the  result.    The  formation  with  silver 
nitrate  of  a  white  precipitate  which  is  soluble  in  ammonium 
hydroxide  and  insoluble  in  nitric  acid  serves  as  a  good  test 
for  chlorides.    Will  this  test  also  serve  to  detect  the  presence 
of  hydrochloric  acid  ? 

[75] 


101.  Weight  of  i  liter  of  hydrogen  chloride.    The  weight  of 
1  liter  of  hydrogen  chloride  may  be  determined  by  the  method 
employed  in  determining  the   weight  of   1   liter  of   carbon 
dioxide  (§52).    Prepare  the  gas  as  described  in  §  64.    The 
gas  should  be  bubbled  through  concentrated  sulfuric  acid  in 
bottle  B  (Fig.  34),  in  order  to  dry  it. 

102.  Preparation  and  properties  of  bromine.    (The  vapor  of 
bromine  must  not  be  inhaled.     Perform  the  experiments  in  a 
hood.)    a.  Use  the  apparatus  employed  in  the  preparation  of 
nitric  acid  (Fig.  40).    Put  into  the  retort  a  mixture  of  2  g.  of 
potassium  bromide  or  sodium  bromide  and  4  g.  of  manganese 
dioxide,  and  add  to  this  through  a  funnel  a  cold  dilute  solu- 
tion of  sulfuric  acid,  prepared  by  slowly  adding  5  cc.  of  sul- 
furic acid  to  10  cc.  of  water.    Shake  the  retort  so  as  to  mix 
the  contents  thoroughly.    The  test-tube  receiver  should  con- 
tain sufficient  water  to  allow  the  end  of  the  retort  to  dip  just 
below  its  surface.    Now  heat  the  retort  gently.    The  bromine 
is  liberated  and  distills  over  (R).    Continue  the  heating  until 
all  the  bromine  has  been  collected. 

6.  Note  the  properties  of  the  bromine  collected  in  the 
bottom  of  the  receiver.  Has  the  water  dissolved  any  of  it  ? 
What  property  is  implied  in  the  name  of  the  element  ?  Test 
the  bleaching  conduct  of  the  aqueous  solution.  How  does  it 
compare  with  chlorine  as  a  bleaching  agent  ?  Try  the  effect 
of  a  little  of  the  bromine  water  on  starch  solution. 

c.  Pour  2  or  3  cc.  of  the  solution  of  bromine  into  each  of 
two  test  tubes.  To  the  one  add  an  equal  volume  of  carbon 
disulfide  and  shake  it  vigorously.  What  is  the  distribution  of 
the  bromine  between  the  two  solvents  ?  To  the  other  test 
tube  add  carbon  tetrachloride  or  chloroform  and  compare  the 
results  with  those  obtained  with  carbon  disulfide. 

103.  Hydrobromic  acid  and  the  bromides,    a.  Add  a  few 
drops  of  sulfuric  acid  to  some  crystals  of  potassium  bromide. 
Note  the  formation  of  white  fumes,  and  also  of  a  reddish- 
brown  vapor.    Give  the  composition  of  each  and  account  for 
its  formation.    (Recall  that  concentrated  sulfuric  acid  is  an 
oxidizing  agent  and  hydrobromic  acid  is  a  reducing  agent.) 

[76] 


b.  Dissolve  in  water  a  crystal  of  potassium  bromide  or  of 
sodium  bromide  and  apply  the  silver  nitrate  test  for  chlorides 
(§  100).    Describe  the  results  (R). 

c.  To  a  solution  of  a  crystal  of  a  bromide  add  a  little 
chlorine  water.    How  do  you  account  for  the  change  in  color  ? 
Shake  the  solution  with  2  or  3  cc.  of  carbon  tetrachloride. 
Does  this  make  the  test  more  delicate  ? 

104.  Preparation  and  properties  of  iodine.    In  a  large  test 
tube  place  a  mixture  of  2  g.  of  potassium  iodide  and  4  g.  of 
manganese  dioxide.    Pour  over  this  mixture  5  cc.  of  sulfuric 
acid,  place  an  empty  funnel  in  the  tube  to  serve  as  a  loose 
stopper,  and  apply  a  gentle  heat  (R).    Note  the  vapor  of  the 
iodine  and  the  grayish-black  crystals  which  are  soon  deposited 
on  the  sides  of  the  test  tube.    What  property  does  the  name 
of  the  element  suggest  ?    Half  fill  two  test  tubes  with  starch 
solution.    Add  to  the  first  a  few  drops  of  a  solution  of  iodine 
prepared  by  shaking  a  small  crystal  in  water.    To  the  second 
add  a  few  drops  of  an  aqueous  solution  of  potassium  iodide. 
Xote  the  results.    Now  add  to  the  second  tube  a  little  chlo- 
rine water  (R).    Does  the  chlorine  water  alone  change  the 
starch  ?    Explain  the  results.    Add   about   2  cc.   of  carbon 
tetrachloride,  and  shake  the  mixture.    In  which  solvent  is  the 
iodine  more  soluble  ?    Dissolve  a  crystal  of  iodine  in  alcohol. 
This  solution  is  called  tincture  of  iodine. 

105.  Hydriodic  acid   and  the   iodides,    a.  Repeat  the   ex- 
periment of  §  103,  a  and  &,  substituting  potassium  iodide  for 
the  bromide.    Note  and  explain  the  results. 

b.  Grind  to  a  fine  powder  two  or  three  crystals  of  iodine 
in  a  mortar  and  add  a  little  water.  Transfer  the  resulting 
mixture  to  a  test  tube,  adding  sufficient  water  to  nearly  fill 
the  tube.  Now  pass  a  slow  current  of  hydrogen  sulfide 
through  the  liquid  until  the  iodine  all  disappears.  Filter  off 
the  white  solid  (explain)  and  boil  the  filtrate  until  the  hydro- 
gen sulfide  is  expelled  (to  test  for  the  presence  of  hydrogen 
sulfide,  hold  in  the  vapor  a  strip  of  paper  moistened  with  a 
solution  of  a  salt  of  lead ;  if  the  acid  is  present  in  the  vapor, 
the  strip  turns  dark,  owing  to  the  formation  of  black  lead 

[77] 


sulfide).    Add  a  few  drops  of  a  solution  of  silver  nitrate  to 
the  resulting  liquid.    Explain. 

106.  To  distinguish  between  chlorides,  bromides,  and  iodides. 
Can  you  distinguish  between  chlorides,  bromides,  and  iodides 
by  means  of  the  silver  nitrate  test  ?    Try  the  following  test 
and  describe  the  results:   Place  in  separate  test  tubes  about 
5  cc.  of  solutions  of  each  of  the  following  compounds:  sodium 
chloride,  sodium  bromide,  sodium  iodide.    Add  to  each  about 
1  cc.  of  carbon  disulfide  and  a  few  drops  of  chlorine  water. 
Shake  the  solutions   in  each  test  tube  so  as  to  mix   them 
thoroughly.    If  both  a  bromide  and  an  iodide  were  present, 
could  you  detect  the  bromide  ?    Try  the  addition  of  more 
chlorine  water  to  the  iodide.    What  becomes  of  the  iodine 
(remember  that  the  chlorine  is  a  good  oxidizing  agent  in  the 
presence  of  water)  ? 

107.  Hypochlorites  and  chlorates,    a.  Dissolve  about  6  g.  of 
solid  potassium  hydroxide  in  18  cc.  of  cold  water,  and  pour 
half  of  the  solution  into  a  good-sized  test  tube,  reserving  the 
other  half  for  §  107,  c.    Fit  up  a  small  chlorine  generator  and 
pass  a  rapid  current  of  chlorine  into  the  solution,  taking  care 
to  keep  the  test  tube  cold  by  immersing  it  in  cold  water. 
When  no  more  chlorine  is  absorbed  (ten  to  fifteen  minutes) 
and  the  solution  is  no  longer  strongly  soapy  to  the  touch, 
withdraw  the  test  tube  and  replace  it  by  the  one  containing 
the  other  half  of  the  solution  of  potassium  hydroxide.    This 
test  tube  should  be  previously  heated  until  it  is  moderately 
warm,  and  it  should  not  be  cooled  while  the  chlorine  is  passed 
in.    Continue  the  operation  until  the  solution  is  saturated. 

&.  Pour  the  contents  of  the  first  test  tube  into  a  beaker. 
To  1  cc.  of  the  solution  add  an  excess  of  nitric  acid  and  then 
a  few  drops  of  silver  nitrate.  What  inference  do  you  draw  ? 
Place  a  piece  of  brightly  colored  cloth  in  the  remainder  of 
the  solution.  Is  it  bleached  ?  Now  add  an  excess  of  dilute 
sulf uric  acid.  Is  the  bleaching  more  rapid  than  before  ?  How 
does  the  sulfuric  acid  assist  in  the  bleaching? 

c.  Pour  the  contents  of  the  second  test  tube  into  a  beaker 
and  set  aside  to  cool,  if  necessary,  till  the  next  period.  Filter 

[78] 


off  the  crystals  and  wash  them  with  a  little  cold  water,  then 
dry  them  on  filter  paper.  Dissolve  a  crystal  in  water  and  test 
with  silver  nitrate  (R).  Is  the  substance  a  chloride  ?  Heat 
the  crystals  in  a  test  tube  and  test  for  oxygen.  Dissolve  the 
residue  in  water  and  test  for  the  presence  of  a  chloride.  Add 
a  little  silver  nitrate  to  the  solution  from  which  the  crystals 
were  obtained.  Is  a  chloride  present?  Write  the  equations 
for  the  action  of  chlorine  on  potassium  hydroxide. 

108.  Bleaching-powder.  Arrange  an  apparatus  according 
to  Fig.  44.  In  this  apparatus  A  represents  a  250-cc.  flask  in 
which  chlorine  is 
generated  accord- 
ing to  §  62,  d, 
using  15g.  of  man- 
ganese dioxide  and 
75  cc.  of  hydro- 
chloric acid.  The 
chlorine  must  be 
evolved  slowly,  and 
so  only  a  very  gentle 
heat  is  applied  to 
the  flask.  The  gas  is 
conducted  through  FlG  44 

a  little  water  in  B, 

then  into  the  tube  (7,  which  is  half  filled  with  slaked  lime. 
The  bottle  D  contains  a  solution  of  sodium  hydroxide.  When 
all  the  chlorine  has  passed  over  from  A,  disconnect  the  appa- 
ratus and  transfer  the  contents  of  the  tube  C  to  a  beaker 
(200-cc.  to  300-cc.)  and  pour  over  it  a  little  sulfuric  acid 
diluted  with  an  equal  volume  of  water.  Hang  a  moist  strip 
of  red  calico  in  the  beaker  and  cover  the  beaker  with  a  glass 
plate.  Note  the  results.  Complete  the  following  equations: 

Ca(OH)2  +  Cl2= 


What  is  the  function  of  the  liquids  in  B  and  D? 

[79] 


CHAPTER  XVIII 


SOME  COMPOUNDS  OF  CARBON 

NOTE.  Students  will  select  either  §  109  or  §  110,  according  to  whether 
formic  acid  or  oxalic  acid  is  to  be  used  in  the  preparation  of  carbon 
monoxide.  It  must  be  kept  in  mind  that  carbon  monoxide  is  a  very  poisonous 
gas.  It  must  not  be  allowed  to  escape  in  the  laboratory. 

109.  Carbon  monoxide :  its  preparation  from  formic  acid. 
a.  Arrange  an  apparatus  according  to  Fig.  45.  Remove  the 
stopper  from  the  flask  B 
(Fig.  45),  pour  in  15  cc. 
of  sulfuric  acid  and 
connect  the  apparatus 
as  shown  in  the  figure. 
Close  the  clamp  C  and 
partially  fill  the  funnel^ 
with  formic  acid  (50 
per  cent  acid).  Now 
open  the  clamp  care- 
fully so  that  the  formic 
acid  will  enter  the  flask,  FIG.  45 

a  drop  at  a  time.  Allow 

8  or  10  drops  to  flow  in  ;  then  close  the  clamp.  If  the  reaction 
does  not  begin  (as  indicated  by  absence  of  effervescence  of 
the  liquid  in  the  flask  and  escape  of  gas  through  the  exit 
tube),  heat  the  flask  very  gently  until  the  reaction  starts ; 
then  open  the  clamp  again  and  admit  the  formic  acid,  a  drop 
at  a  time,  so  as  to  secure  a  regular  flow  of  gas  from  the  flask. 
If  necessary,  add  more  formic  acid  to  the  funnel  so  as  to  keep 
it  partially  filled  (?).  Collect  three  bottles  of  the  gas  as 
shown  in  the  figure.  Close  the  clamp  so  as  to  stop  further 

[80] 


generation  of  gas.  Slip  the  glass  plates  over  the  mouths  of 
the  bottles  and  remove  the  bottles  from  the  trough.  In  the 
first  bottle  filled,  is  the  gas  pure  carbon  monoxide  ?  Remove 
the  glass  cover  and  test  it  with  a  flame  (?).  Repeat  with  the 
second  bottle  filled  (?).  Slip  the  glass  plate  from  the  third 
bottle  just  far  enough  to  pour  into  the  bottle  5  cc.  of  clear 
limewater ;  then  quickly  replace  the  glass  plate  and,  holding 
it  firmly  against  the  mouth  of  the  bottle,  shake  the  contents 
of  the  bottle.  Note  any  change  in  the  appearance  of  the  lime- 
water.  Now  tip  the  bottle  as  far  as  possible  without  spilling 
the  limewater;  remove  the 

A /? 

glass  plate  and  quickly  ignite 
the  gas,  holding  the  bottle  in 
this  position  so  that  at  least 
a  portion  of  the  combustion 


I 

0 


product  may  be  retained  in  FlG  46 

the  bottle.    When  the  flame 

dies  out,  at  once  cover  the  mouth  of  the  bottle  with  the  glass 

plate  and  shake  the  contents.    Note  the  results. 

b.  Introduce  into  the  tube  A  (Fig.  46),  a  small  amount  of 
copper  oxide  and  arrange  the  apparatus  as  shown  in  the 
figure.  Now  connect  the  exit  tube  D  (Fig.  45)  with  the  tube 
A  (Fig.  46).  Heat  the  copper  oxide  gently ;  at  the  same 
time  pass  a  slow  current  of  carbon  monoxide  through  the 
tube,  generating  the  gas  as  described  in  a.  Continue  until 
the  limewater  and  copper  oxide  both  have  visibly  changed. 
Describe  the  results  and  write  the  equations  for  all  the 
reactions  involved. 

110.  Carbon  monoxide  :  its  preparation  from  oxalic  acid. 
a.  Arrange  an  apparatus  according  to  Fig.  47.  The  flask  A 
should  have  a  capacity  of  from  100  to  150  cc.  The  bottle  B 
contains  a  solution  of  sodium  hydroxide,  and  C  and  D  con- 
tain solutions  of  limewater.  The  hard-glass  tube  E  (12  to 
15  cm.  in  length)  contains,  near  its  middle,  a  layer  of  1  or 
2  g.  of  black  copper  oxide.  The  whole  apparatus  must  be 
air-tight.  After  the  apparatus  has  been  approved  by  the  in- 
structor, put  7  or  8  g.  of  oxalic  acid  (H2C2O4)  into  the  flask  A, 

[81] 


and  pour  over  this  25  cc.  of  concentrated  sulfuric  acid.  Stopper 
the  flask  tightly  and  apply  a  very  gentle  heat,  at  the  same 
time  heating  the  copper  oxide  in  E.  Regulate  the  heat  so  as 
to  cause  the  evolved  gas  to  bubble  slowly  through  B  and  C. 
Collect  over  water  in  F  any  gas  escaping  from  D.  What 
evidences  have  you  that  the  copper  oxide  is  reduced? 

b.  Disconnect  the  bottle  B  and  attach  a  rubber  delivery 
tube  in  its  place.  Collect  by  displacement  of  water  a  test  tube 
full  of  the  gas  evolved  in  A.  Close  the  mouth  of  the  tube 


FIG.  47 

with  the  thumb  and  invert  it  in  a  beaker  containing  a  solu- 
tion of  sodium  hydroxide.  What  proportion  of  the  gas  is 
absorbed?  What  is  it?  What  gas  remains? 

c.  Withdraw  the  heat  from  A,  remove  the  stoppers  from 
A,  B,  and  (7,  and  test  with  a  burning  splint  the  gas  inclosed 
in  each.  Account  for  the  results.  Also  test  with  a  burning 
splint  the  gas  which  escaped  from  1>,  and  account  for  the 
result.  Why  is  the  gas  passed  through  a  solution  of  sodium 
hydroxide?  Write  the  equations  for  the  reactions  which  take 
place  in  each  of  the  containers  A,  B,  C,  D. 

111.  Carbonic  acid  and  its  salts,  a.  Generate  carbon 
dioxide  as  in  §  51.  Wash  the  gas  by  bubbling  it  through 
water,  and  then  pass  it  through  pure  water  in  a  beaker  until 
the  gas  is  no  longer  absorbed.  Test  the  solution  with  litmus. 
How  does  the  taste  of  the  solution  compare  with  that  of 
water  ? 

[82] 


b.  Give  the  formula  and  properties  of  the  acid  of  which 
the  carbonates  are  salts.    Try  the  action  of  hydrochloric  acid 
on  a  small  amount   (about  1  g.)   of   each   of   the  following 
carbonates :   sodium  carbonate,  potassium  carbonate,  magne- 
sium carbonate  (R).    How  can  you  detect  the  presence  of 
carbonates  ?    Is  limestone  a  carbonate  ?    Is  the  action  of  sul- 
furic  and  nitric  acids  on  carbonates  similar  to  that  of  hydro- 
chloric acid  (R)  ?    Why  is  carbonic  acid  so  readily  liberated 
from  carbonates  ? 

c.  Pass  a  current  of  carbon  dioxide  through  10  cc.  of  the 
laboratory  solution  of  sodium  hydroxide  (1  part  of  sodium 
hydroxide  to  10  parts  of  water).    Does  the  solution  remain 
clear  ?    When  no  more  of  the  gas  is  absorbed,  evaporate  the 
solution  to  dryness  and  test  the  residue  for  carbonates.    What 
is  the  residue  (R)  ? 

d.  In  what  respects  does  carbon  dioxide  resemble  sulfur 
dioxide  ? 

e.  Half  fill  a  small  beaker  with  limewater  and  pass  carbon 
dioxide  through  the  liquid  (R).    Continue  until  the  precipi- 
tate which  at  first  forms  is  dissolved  (R),  and  then  divide  the 
solution  into  two  parts.    To  the  one  add  a  little  clear  lime- 
water  (R)  ;  to  the  other  apply  heat  until  it  boils  rapidly  (R). 
Account  for  the  fact  that  carbon  dioxide  will  cause  a  precipi- 
tate in  a  solution  of  calcium  hydroxide  but  not  in  a  solution 
of  sodium  hydroxide. 

112.  Oxidation  of  urea.  Prepare  about  100  cc.  of  a  solution 
of  sodium  hypochlorite  by  passing  chlorine  into  a  cold  solu- 
tion of  sodium  hydroxide  (§  107)  until  the  chlorine  is  no 
longer  readily  absorbed,  thus  leaving  some  sodium  hydroxide 
present  in  solution  with  the  sodium  hypochlorite  (R).  Fill  a 
test  tube  with  the  solution  and  invert  it  in  a  small  evaporat- 
ing dish  partially  filled  with  the  same  solution.  Now,  by 
means  of  a  medicine  dropper,  the  small  end  of  which  is  curved 
slightly,  introduce  into  the  solution  in  the  test  tube  a  few 
drops  of  a  solution  of  urea  (R). 


[83] 


CHAPTER  XIX' 


THE  LAW  OF  GAY-LUSSAC 

NOTE.    Before  performing  the  experiment,  read  the  directions  care- 
fully and  report  to  the  instructor  for  quiz  on  the  methods. 

113.  Determination  of  the  volume  of  hydrogen  obtained 
from  a  known  volume  of  hydrogen  chloride.  Prepare  some 
sodium  amalgam  as  follows :  Pour  about  10  cc.  of  mercury 
into  an  evaporating-dish  and  heat  slightly  (hood).  Add,  one 
at  a  time,  five  pieces  of  sodium,  each 
as  large  as  a  pea.  If  each  bit  of  the 
sodium  does  not  combine  with  the 
mercury  on  coming  in  contact  with  it, 
start  the  combination  by  pushing  the 
sodium  under  the  surface  of  the  mer- 
cury by  means  of  a  long  glass  rod. 
After  the  sodium  has  all  been  added, 
pour  the  resulting  amalgam  into  a  small 
bottle  and  stopper  it. 

Obtain  from  the  storeroom  a  tube 
about  50  cm.  in  length  and  15  mm.  in 
diameter  (Fig.  48,  (7),  and  fill  it  with  dry  hydrogen 
chloride.    To  do  this,  arrange  an  apparatus  according 
to  Fig.  48.    Generate  hydrogen  chloride  in  At  dry  it 
by  passing  it  through  the  concentrated  sulfuric  acid 
in  B,  and  conduct  it  into  C  by  means  of  a  long  deliv- 
ery tube.  When  the  tube  C  is  completely  filled  with  the 
gas,  slowly  withdraw  it  from  the  delivery  tube,  pour 
in  the  sodium  amalgam  prepared  above,  and  at  once  close  the 
mouth  of  the  tube  firmly  with  the  thumb,  slightly  moistened. 
Now  shake  the  tube  so  as  to  bring  the  gas  into  contact  with 

[84] 


FIG.  48 


the  amalgam.  Finally,  invert  the  tube  in  a  vessel  of  water 
and  remove  the  thumb.  The  water  rises  in  the  tube.  Deter- 
mine the  nature  and  the  exact  volume  of  the  resulting 
gas.  Draw  conclusions  with  reference  to  the  composition  of 
.hydrogen  chloride. 

114.  Determination  of  the  relative  volumes  of  nitrogen  and 
hydrogen  obtained  by  the  decomposition  of  ammonia.  Use  the 
apparatus  shown  in  Fig.  32,  except  that  the  glass  tube  C 
used  in  §  113  is  substituted  for  the  test  tube.  Clean  and 
dry  the  tube  C  and  fill  it  with  dry  chlorine  gas  (hood).  The 
chlorine  is  generated  in  A  (Fig.  48),  is  dried  by  passing  it 
through  the  sulfuric  acid  in  B,  and  is  conducted  into  C  by 
means  of  the  long  delivery  tube.  While  the  tube  is  being 
filled,  pour  into  the  funnel  A  (Fig.  32)  a  few  cubic  centi- 
meters of  a  concentrated  solution  of  ammonium  hydroxide. 
Open  the  clamp  B  until  the  rubber  tube  and  the  glass  nozzle 
are  completely  filled  with  the  solution ;  then  close  the  clamp. 
When  the  tube  C  is  completely  filled  with  chlorine,  slowly 
withdraw  it  from  the  delivery  tube  and  at  once  connect  it 
with  B  (Fig.  32)  by  means  of  the  cork  (the  cork  must  be 
firmly  inserted  so  as  to  make  the  connection  air-tight).  Now 
allow  the  ammonium  hydroxide  solution  to  enter  the  tube,  a 
drop  at  a  time.  The  chlorine  combines  with  the  hydrogen  of 
the  ammonia  and  liberates  the  nitrogen  (R).  When  the  addi- 
tion of  the  hydroxide  solution  no  longer  causes  any  action, 
pour  into  the  funnel  some  dilute  sulfuric  acid  and.  admit  this 
to  the  tube  so  as  to  combine  with  the  excess  of  ammonia.  The 
tube  should  be  shaken  so  as  to  mix  the  contents  thoroughly. 
When  the  sulfuric  acid  no  longer  causes  any  action,  fill  the 
funnel  with  water  and  open  the  screw  clamp  (care  must  be 
taken  that  no  air  is  admitted  to  the  tube  with  the  liquids). 
When  no  more  water  will  enter  the  tube,  close  the  screw 
clamp,  remove  the  funnel,  and  invert  the  tube  in  a  trough  of 
water.  Remove  the  cork,  adjust  the  pressure,  and  determine 
the  exact  volume  and  the  nature  of  the  remaining  gas. 
What  volume  of  hydrogen  was  necessary  to  combine  with 
the  volume  of  chlorine  taken  ?  What  was  the  source  of  this 

[85] 


hydrogen  ?  What  became  of  the  liberated  nitrogen  ?  How 
does  the  volume  of  nitrogen  liberated  compare  with  the 
volume  of  hydrogen  used  in  combining  with  the  chlorine  ? 
What  are  the  relative  volumes  of  hydrogen  and  nitrogen  set 
free  by  the  decomposition  of  ammonia  ? 

115.  Determination  of  the  volume  of  nitrogen  obtained  from 
a  known  volume  of  ammonia,    a.  Use  the  same  apparatus  as 
in  §  114.    Clean  and  dry  the  tube   C  and  fill  it  with  dry 
ammonia  gas  by  upward  displacement.    The  gas  is  most  con- 
veniently generated  by  warming  a  concentrated  solution  of 
ammonium  hydroxide.    It  is  dried  by  passing  it  through  a 
tube  filled  with  small  pieces  of  lime,  and  is  then  conducted 
through  the  long  delivery  tube  into  C.    While  the  tube  is 
being  filled,  pour  into  the  funnel  (Fig.  32)    a  solution  of 
sodium  hypobromite  (side  shelf).    Open  the  clamp  until  the 
rubber  tube  and  the  nozzle  are  completely  filled  with  the  solu- 
tion, then  close  the  clamp.    When  the  tube  C  is  completely 
filled  with  ammonia,  slowly  withdraw  the  delivery  tube  and 
quickly  connect   (air-tight)  the  tube  with  B   (Fig.  32)  by 
means  of  the  cork.     Now  admit  the  hypobromite  solution 
slowly  as  long  as  any  action  takes  place.     The  following 
equation  represents  the  reaction  involved: 

3  NaBrO  +  2  NH3  =  3  NaBr  +  3  H2O  +  N2 

When  the  action  has  ceased,  water  is  admitted.  When  no 
more  water  will  enter  the  tube,  remove  the  funnel,  invert 
the  tube  in  a  trough  of  water,  and  remove  the  cork.  Adjust 
the  pressure  and  determine  the  volume  and  nature  of  the 
remaining  gas.  How  does  the  volume  of  nitrogen  obtained 
from  the  ammonia  compare  with  the  volume  of  the  ammonia? 
b.  From  the  results  obtained  in  §§  114  and  115,  c,  state 
the  relations  between  the  volumes  of  hydrogen  and  nitrogen 
which  combine  to  form  ammonia,  also  their  relation  to  the 
volume  of  ammonia  formed. 

116.  Comparison  of  results  obtained.     Recall  (lecture  ex- 
periment) the  relation  between  the  volumes  of  hydrogen  and 
oxygen  which  combine  to  form  water,  also  the  volume  of 

[86] 


water  vapor  formed.  Recall  also  the  relation  between  the 
volumes  of  hydrogen  and  chlorine  which  combine  to  form 
hydrogen  chlorine,  and  the  relation  each  holds  to  the  volume 
of  hydrogen  chloride  formed.  Represent  graphically  the  pro- 
portion by  volume  in  which  the  two  elements  in  each  of  the 
following  pairs  of  elements  combine,  and  also  the  relative 
volume  of  the  compound  formed  in  each  case  by  their  union : 
hydrogen  and  chlorine ;  hydrogen  and  oxygen ;  hydrogen  and 
nitrogen.  State  all  of  these  facts  in  the  form  of  a  generali- 
zation. What  is  this  generalization  called? 


[87] 


CHAPTER  XX 
COMBINING  WEIGHTS  AND  MOLECULAR  WEIGHTS 

117.  The   combining   weight  of   zinc.    From   your   results 
obtained  in  §  49  calculate  the  combining  weight  of  zinc  re- 
ferred to  hydrogen  as  unity.    Repeat  the  experiment,  substi- 
tuting for  the  sulfuric  acid  a  solution  of  hydrochloric  acid, 
prepared  by  adding  1  volume  of  the  laboratory  acid  (density 
1.12)  to  1  volume  of  water.    Compare  the  results. 

118.  The  combining  weight  of  magnesium.    Determine  the 
combining  weight    of   magnesium,  referred   to    hydrogen   as 
unity,  by  dissolving  the  metal  in  dilute  hydrochloric  acid  and 
measuring  the  hydrogen  evolved.    About  0.5  g.  of  magnesium 
ribbon,  prepared  as  directed  in  §  44,  should  be  used.  The  hydro- 
chloric acid  is  prepared  by  adding  1  volume  of  the  laboratory 
acid  (density  1.12)  to  4  volumes  of  water.    The  apparatus 
used  is  shown  in  Fig.  27. 

119.  The  combining  weight  of  aluminium.   Follow  the  same 
method  as  was  used  in  the  determination  of  the  combining 
weights  of  zinc  and  magnesium.   The  hydrochloric  acid  used 
is  prepared  by  diluting  2  volumes  of  the  laboratory  acid  with 
1  volume  of  water. 

120.  Determination  of  molecular  weights.    Molecular  weight 
of  chloroform.    (The  following  experiment  will  be  performed 
by  the  laboratory  instructor.    The  students  will  answer  all 
questions  and  make  the  calculations.    Before  performing  the 
experiment,  read  over  the  directions  and  discuss  the  method 
until  you  thoroughly  understand  it.)    Obtain  from  the  store- 
room the  Victor  Meyer  apparatus   shown  in  Fig.  49.     This 
apparatus   consists   of  two  glass  tubes,  the   smaller  one   of 

[88] 


E 


D 


which  is  suspended  in  the  larger,  as  represented  in  the  dia- 
gram. The  liquid  in  the  outer  tube  D  is  water.  In  the 
bottom  of  the  inner  tube  A  is  placed  a  little  asbestos  fiber 
or  sand.  The  graduated  tube  E  has  a  capacity  of  about  50  cc. 
Arrange  the  apparatus  as  shown  in  the  figure,  except  that 
the  tube  E  is  not  brought  over  the  exit  tube  C.  Tightly 
stopper  the  tube  A  with  a  rubber  stopper  provided  with  a 
stopcock  B,  and  heat  the  water  in  D  to  boiling ;  then  regulate 
the  flame  so  that  a  uniform  heat  is  applied, 
the  heat  being  sufficient  to  keep  the  water 
boiling.  The  air  in  tube  A  is  heated  by 
the  steam,  and  some  of  it  escapes  through 
Cj  bubbling  up  through  the  water.  While 
the  heat  is  being  applied,  wrap  the  end  of 
a  fine  platinum  wire  5  or  6  cm.  in  length 
about  the  neck  of  the  little  glass-stoppered 
bottle  F,  suspend  the  bottle  from  the  beam 
of  the  balance,  and  weigh  it.  Completely 
fill  the  bottle  with  chloroform,  insert  the 
stopper,  carefully  wipe  the  bottle  to  remove 
any  of  the  liquid  adhering  to  the  outside, 
and  reweigh.  Now  observe  whether  the 
air  is  still  escaping  from  tube  C.  If  not, 
open  the  stopcock  B,  remove  the  stopper 
and  stopcock  j5,  and,  after  loosening  the 
stopper  in  the  little  glass  bottle,  hold  the 
bottle  in  a  vertical  position  with  the  plat- 
inum wire  and  drop  it  into  the  tube  A. 
As  quickly  as  possible  insert  the  stopper  * 
and  close  B.  Before  any  gas  can  escape, 
bring  the  tube  E  (which  must  be  completely  filled  with 
water)  over  the  end  of  the  tube  C  and  clamp  it  in  this 
position.  The  bottle  F  drops  to  the  bottom  of  the  tube,  and 
the  chloroform,  on  account  of  its  low  boiling  point  (61°),  is 
vaporized  by  the  heat.  The  vapor  formed  expels  a  definite 
volume  of  air,  which  is  caught  in  the  graduated  tube  E. 
The  heating  is  continued  until  the  gas  ceases  to  escape. 

[89] 


FIG.  49 


The    stopcock  B  is  then   opened    arid  the    heat  withdrawn. 
Insert  the  values  in  the  f  ollowm£  table : 


Weight  of  chloroform  taken 

Volume  of  air  collected  in  E 

Temperature  of  air  in  E 

Pressure  to  which  the  air  in  E  is  subjected  . 

It  is  evident  that  the  volume  of  the  air  in  E  must  be  equal 
to  the  volume  occupied  by  the  vapor  formed  by  the  chloro- 
form, provided  that  the  vapor  is  measured  under  the  condi- 
tions of  temperature  and  pressure  to  which  the  air  in  the  tube 
is  subjected.  Knowing,  therefore,  the  weight  of  the  chloro- 
form taken  and  the  volume  of  vapor  formed  under  definite 
conditions,  it  is  easy  to  calculate  the  weight  of  22.4  liters  of 
the  vapor  under  standard  conditions. 

By  analysis  it  is  possible  to  show  that  chloroform  has  the 
following  composition :  carbon,  10.04  per  cent;  hydrogen,  0.84 
per  cent;  chlorine,  89.12  per  cent.  From  your  results  deter- 
mine the  formula  of  chloroform.  Why  is  it  necessary  to  know 
the  molecular  weight  in  order  to  determine  the  formula? 

Notes  and  queries.  The  stopper  of  the  little  bottle  F  con- 
taining the  chloroform  is  loosened  before  dropping  the  bottle 
into  the  tube,  in  order  that  the  vapor  of  the  chloroform  may 
easily  escape.  In  performing  the  experiment  the  heat  must 
be  applied  uniformly  until  the  experiment  is  completed 
(why?);  hence  the  flame  should  be  protected  from  drafts 
of  air.  The  amount  of  chloroform  introduced  must  be  small 
(about  0.13  g.)  ;  otherwise  the  volume  of  the  vapor  will  be 
so  large  that  the  tube  E  will  not  hold  the  displaced  air. 
Could  other  liquids  be  substituted  for  water  in  the  outer  tube? 
Before  performing  the  experiment  the  tube  A  must  be  per- 
fectly dry  (why  ?).  To  remove  any  vapor,  insert  a  long, 
narrow  glass  tube  and  force  a  current  of  dry  air  through 
the  tube,  at  the  same  time  applying  a  gentle  heat.  Before 
measuring  the  volume  of  the  air  in  E,  bring  the  water  within 
the  tube  to  the  level  of  the  water  outside,  if  necessary  trans- 
ferring the  tube  to  a  deep  cylinder  to  accomplish  this. 

[90] 


CHAPTER  XXI 
SOME  HYDROCARBONS 

121.  Methane,    a.  Weigh  out  approximately  10  g.  of  fused 
sodium  acetate  and  double  that  weight  of  soda  lime  (a  mix- 
ture of  calcium  oxide  and  sodium  hydroxide).    Grind  the  two 
together  in  a  mortar  and  put  the  mixture  in  a  hard-glass  test 
tube  furnished  with  a  stopper  and  a  delivery  tube  (the  oxygen 
generator  of  §  15).    Support  the  tube  with  a  clamp  and  heat 
gently,  gradually  increasing  the  heat.    After  the  air  has  been 
expelled  from  the  apparatus,  collect  two  or  three  bottles  of 
the  gas  over  water. 

&.  Note  the  color,  odor,  and  solubility  of  the  gas.  Hold  a 
lighted  match  to  the  bubbles  as  they  escape  from  the  water. 
Pour  a  little  bromine  water  into  one  of  the  bottles,  place  a 
stopper  in  the  bottle,  and  shake  vigorously.  Is  the  color  of 
the  bromine  changed?  Set  fire  to  the  gas  in  the  second  bottle 
and,  after  it  has  burned,  test  the  residual  gas  for  the  presence 
of  carbon  dioxide.  Half  fill  a  bottle  with  methane  and  then 
displace  the  remaining  water  with  air.  Quickly  bring  the 
mouth  of  the  bottle  over  a  flame.  What  connection  is  there 
between  this  experiment  and  a  mine  explosion  ? 

122.  Ethylene.    a.  Arrange  a  250-cc.  generating-flask  on  a 
sand  bath  and  connect  it  with  a  gas- washing  bottle  containing 
dilute  sodium  hydroxide  and  provided  with  a  delivery  tube 
for  collecting  over  water.    Pour  10  cc.  of  alcohol  very  slowly 
into  60  cc.  of  concentrated  sulfuric  acid,  stirring  constantly 
during  the  mixing ;  pour  the   mixture   into  the  generating- 
flask  and  insert  the  stopper.    Carefully  heat  the  mixture  until 
there  is  a  steady  evolution  of  gas.    If  the  heating  is  too  rapid 
or  the  temperature  reached  is  too  high,  the  mixture  may  foam 

[91] 


badly  and  the  alcohol  be  carbonized.  Remembering  that  some 
free  carbon  is  usually  formed  in  the  reaction,  what  gases  other 
than  ethylene  would  you  expect  to  be  formed  ?  How  would 
these  act  with  sodium  hydroxide  in  the  wash  bottle  ? 

b.  Collect  two  bottles  of  the  gas  and  test  its  combusti- 
bility and  its  conduct  towards  bromine  water.  How  does  it 
compare  with  methane  ? 

123.  Acetylene,    a.  With  your  pliers  pick  up  a  small  piece 
of  calcium  carbide  and  quickly  insert  it  under  the  mouth  of 
a  test  tube  filled  with  water  and  inverted  in  a  beaker  of  water. 
What  action  takes  place  ?  Test  the  reaction  of  the  water 
with  litmus  paper.    Holding   the    test  tube   in   an    inverted 
position  (why  ?),  remove  it  from  the  water  and  quickly  ignite 
the  gas  at  the  mouth  of  the  tube.    How  does  the  flame  com- 
pare with  that  of  burning  ethylene  and  burning  methane  ? 

b.  Collect  another  test  tube  full  of  the  gas,  quickly  pour 
into  it  about  1  cc.  of  bromine  water,  close  the  mouth  of  the 
tube  with  the  thumb,  and  shake  it  vigorously.  Is  the  bromine 
absorbed  by  the  acetylene  ? 

124.  lodoform.    Pour  2  or  3  cc.  of  a  solution  of  sodium  hy- 
droxide into  a  test  tube,  add  a  crystal  or  two  of  iodine,  and 
then  a  few  drops  of  alcohol.    Warm  slightly  and  observe  the 
changes  which  occur.    What  color  does  the  solution  become  ? 
Is  there  a  precipitate  ?  What  do  you  notice  in  regard  to  the 
odor  ?  The  product  is  iodoform   (CHI3),  and  its  formation 
may  be  used  as  a  good  test  for  alcohol  (though  some  other 
substances,  when  treated  with  iodine  and  an  alkali,  also  give 
iodoform). 


[92] 


CHAPTER  XXII 
FUEL  GASES ;  FLAMES ;  MEASUREMENT  OF  HEAT 

125.  Nature  of  a  flame.    Prepare  some  charcoal  by  heating 
pieces  of  splints  3  or  4  cm.  in  length  in  the  bottom  of  a  test 
tube.    Note  and  account  for  the  difference  between  the  com- 
bustion of  the  splint  and  that  of  the  charcoal.    What  are  the 
conditions  necessary  for  the  production  of  a  flame  ?    Light  a 
candle  and  place  it  so  that  the  flame  is  against  a  black  back- 
ground and  is  not  disturbed   by  air  drafts ;   then  note  the 
different  cones  in  the  flame.    Test  the  relative  temperatures 
of  different  parts  of  the  flame  by  means  of  narrow  strips  of 
splints.    Draw  a  diagram  showing  the  different  parts  of  the 
flame.    Extinguish  the  candle  flame  and  hold  a  lighted  splint 
2  or  3  cm.  from  the  wick  in  the  little  column  of  smoke.    The 
candle  is  relighted.    What  does  the  experiment  prove  ? 

126.  Illuminating  gas.    Repeat  the  experiment  of  §  50,  a, 
substituting  some  pulverized  soft  coal  for  the  bits  of  splint. 
Test  the  inflammability  of  the  gas.    Is  its  flame  luminous  ? 

127.  Products  of  combustion.    What  elements  constitute  the 
main  parts  of  ordinary  fuels?    What  products  form  when 
these  elements  burn  in  air  or  oxygen  ?    Devise  simple  experi- 
ments to  show  the  presence  of  these  products  in  the  gases 
evolved  by  the  burning  candle.    Account  for  the  moisture 
deposited  on  a  lamp  chimney  when  the  lamp  is  first  lighted. 

128.  Kindling  temperature.   What  is  meant  by  the  kindling 
temperature  of  gases  ?    Press  a  piece  of  wire  gauze  halfway 
down  on  a  Bunsen  flame.     Notice  that  the  flame  does  not 
extend  above  the  gauze.    Is  this  due  to  the  absence  there  of 
combustible  gases  ?    Test  for  their  presence  by  means  of  a 
lighted  splint.    Account  for  the  results.    Turn  off  the  gas, 

[93] 


then  turn  it  on  and  ignite  it  over  a  piece  of  wire  gauze  held 
horizontally  4  or  5  cm.  above  the  top  of  the  burner.  Note  the 
results  and  explain.  How  does  the  miner's  safety  lamp  pre- 
vent explosions  ?  Hold  a  porcelain  dish  in  a  small,  luminous 
Bunsen  flame.  Account  for  the  deposition  of  carbon  (carbon 
black).  Does  the  nonluminous  flame  deposit  carbon  ?  To 
what  is  the  luminosity  of  the  flame  due  ? 
i  129.  Bunsen  flame.  Hold  a  wire  horizontally  in  the  base 
of  the  Bunsen  flame  for  two  or  three  seconds  and  note  the  re- 
sults. In  the  same  way  determine  the  relative  temperatures  of 
various  parts  of  the  flame.  Turn  the  gas  down  until  the  flame 
is  7  or  8  cm.  in  height,  then  quickly  thrust  a  piece  of  white 
cardboard,  about  10  cm.  in  height,  vertically  through  the  center 
of  the  flame,  the  lower  edge  of  the  cardboard  resting  against 
the  top  of  the  burner.  Remove  the  card- 
board before  it  is  ignited,  and  note  from 
the  scorched  portions  the  relative  tem- 
peratures of  different  parts  of  the  flame. 
Draw  a  diagram  to  illustrate  your  results. 
That  the  center  of  the  base  of  the 
flame  contains  the  unburned  gas  may  be 
shown  by  holding  in  it  the  end  of  an  in-  FlG  50 

clined  glass  tube  12  or  15  cm.  in  length 
(Fig.  50)  and  igniting  the  gas  at  the  upper  end  of  the  tube. 

130.  Oxidizing-flame  and  reducing-flame.   Ask  the  instructor 
to  show  you  how  to  produce  each  of  these  flames  by  means 
of  the  blowpipe.    Then  heat  in  the  reduciiig-flame  a  small 
amount  of  an  intimate  mixture  of  sodium  carbonate  and  lead 
oxide  (PbO)  placed  in  a  small  cavity  in  a  piece  of  charcoal. 
Have  you  any  evidence  of  the  reduction  of  the  oxide  ?    Try 
the  effect  of  the  oxidizing-flame  on  a  small  piece  of  metallic 
lead  on  charcoal.    Describe  the  results.    Carefully  distinguish 
between  an  oxidizing-flame  and  a  reducing-flame.    To  what 
does  each  owe  its  peculiar  property? 

131.  Heat  of  neutralization.    Prepare  a  calorimeter  as  rep- 
resented in  Fig.  51.    The  calorimeter  beaker  A  should  hold 
approximately  400  cc.,   should  be   of    light  weight,   and,  if 

[94] 


possible,  should  be  without  a  Hp.  The  larger  beaker,  B,  should 
hold  at  least  1  liter.  Provide  three  flat  corks  of  equal  thick- 
ness to  serve  as  supports  f or  A  ;  a  stirrer  C  made  from  a  piece 
of  light-weight  rod  and  bent  into  a  ring  at  the  end ;  and  a 
thermometer  D  graduated  to  one  tenth  of  a  degree.  Cut  a 
cardboard  cover  E,  circular  in  form,  to  fit  over  A  and  inside 
B ;  also  a  cover  F  for  J5,  which  may  be  left  square.  Punch  a 
hole  in  the  middle  of  each  cover,  through  which  the  ther- 
mometer may  be  pushed,  but  small  enough  to  hold  the  latter 
in  place.  A  second  hole  must 
also  be  punched  in  each  cover 
to  provide  for  the  handle  of 
the  stirrer. 

In  a  graduated  cylinder, 
measure  as  accurately  as  pos- 
sible 150  cc.  of  a  normal  solu- 
tion of  NaOH,  and  pour  it 
into  A  ;  measure  also  150  cc.  of 
a  normal  solution  of  HC1,  and 
pour  it  into  a  second  beaker. 
Read  the  temperature  of  the 
two  solutions  at  frequent  in- 
tervals until  they  have  come 
to  exactly  the  same  constant 
point,  warming  the  cooler  one, 
if  necessary,  with  the  hand. 
Make  a  note  of  this  temperature,  reading  it  as  closely  as  pos- 
sible. Arrange  the  thermometer  and  the  stirrer  through  the 
holes  in  the  covers  so  that  all  may  be  put  in  place  together ; 
then  pour  the  acid  into  the  alkali  and  at  once  adjust  the 
covers  and  the  thermometer.  Maintain  a  slow  but  steady 
motion  with  the  stirrer,  watching  the  rise  in  temperature  very 
closely.  Make  a  note  of  the  highest  temperature  reached. 

The  specific  heat  of  the  dilute  solution  may  with  little 
error  be  taken  as  equal  to  that  of  water  (unity),  so  that  the 
rise  in  temperature,  multiplied  by  the  volume  of  the  solution 
(in  this  case  300  cc.),  will  give  the  apparent  heat  of  the 

[95] 


FIG.  61 


reaction.  Some  heat  has  been  absorbed  by  the  calorimeter 
beaker,  the  thermometer,  and  the  stirrer,  and  this  must  be 
determined  and  added  to  the  apparent  heat.  If  the  materials 
are  of  the  character  specified,  experiment  has  shown  that  the 
heat  absorbed  will  be  about  12  cal.,  and  this  value  may  be 
assumed  with  little  error.  The  constant  may  be  calculated 
by  determining  the  weight  of  the  portion  of  the  beaker  and 
stirrer  in  contact  with  the  water,  multiplying  this  by  0.19 
(the  specific  heat  of  glass),  and  to  this  adding  the  volume  of 
the  submerged  part  of  the  thermometer,  multiplied  by  0.49. 

Having  in  this  way  found  the  heat  of  neutralization  of 
150  cc.  of  normal  NaOH,  calculate  the  heat  of  neutralization 
of  1  gram-molecular  weight  of  the  base  (1000  cc.).  The 
value  determined  by  accurate  experiment  is  13,700  cal. 

132.  Heat  of  solution.  Accurately  measure  300  cc.  of  pure 
water  into  the  calorimeter  beaker  A  (Fig.  51).  Grind  about 
15  g.  of  potassium  nitrate  to  a  very  fine  powder  in  a  mortar, 
and  place  it  in  a  test  tube.  Weigh  the  tube  and  nitrate  very 
accurately  and  immerse  the  lower  end  of  the  tube  in  the 
water  in  the  calorimeter  until  the  nitrate  has  come  to  the 
temperature  of  the  water  (about  fifteen  minutes).  Make  a 
note  of  the  temperature.  Remove  the  tube  with  as  little  loss 
of  water  as  possible,  roughly  dry  it  with  a  towel,  and  at  once 
pour  most  of  the  nitrate  into  the  water.  Quickly  replace  the 
covers,  as  in  §  131,  stir  vigorously,  and  note  the  lowest  tem- 
perature recorded.  Then  weigh  the  tube  and  the  remaining 
nitrate,  deducting  this  weight  from  the  original  one  to  get 
the  weight  of  the  nitrate  added.  From  the  data  so  secured, 
together  with  the  calorimeter  constant,  calculate  the  heat 
of  solution  of  1  gram-molecular  weight  of  potassium  nitrate. 
Berthelot  found  this  to  be  -  8300  cal. 


[96] 


CHAPTER  XXIII 
CARBOHYDRATES;   ALCOHOLS;    SOAPS 

133.  Preparation  of  Fehling's  solution.    The  most  common 
test  for  sugars  is  their  reaction  with  the  so-called  "Fehling's 
solution."  This  solution  is  prepared  as  follows :  Dissolve  3.5  g. 
of  copper  sulfate  crystals  in  water  and  dilute  to  50  cc.    Pour 
the  solution  into  a  bottle  and  label  it  "Solution  A."   Dissolve 
5  g.   of  sodium  hydroxide  and  17.5g.  of  sodium-potassium 
tartrate  (Rochelle  salts)  in  about  40  cc.  of  water  and  dilute 
to  50  cc.    Pour  this  into  a  bottle  and  label  it  "  Solution  B." 

134.  Action  of  Fehling's  solution  on  dextrose.    Pour  into  a 
test  tube  about  3  cc.  of  each  of  the  above  solutions  marked 
"A"  and  "  B."    When  thoroughly  mixed,  the  resulting  solu- 
tion should  be  deep  blue  but  perfectly  clear.    Heat  nearly  to 
boiling,  add  a  few  drops  of  a  solution  of  commercial  glucose, 
and  continue  the  heating.    The  copper  compound  in  the  solu- 
tion is  reduced  to  cuprous  oxide  by  the  dextrose  present,  and 
this  separates  in  the  form  of  a  red  or  yellow  solid.    Levulose 
will  act  in  the  same  way.    Test  samples  of  candy,  honey,  and 
molasses  for  the  presence  of  these  sugars. 

135.  Action  of  Fehling's  solution  on  cane  sugar.    In  a  simi- 
lar way  try  the  action  of  cane  sugar  on  the  solution.    Note 
that  the  pure  sugar  does  not  reduce  the  alkaline  copper  com- 
pound.   Now  dissolve  about  1  g.  of  the  sugar  in  10  cc.  of 
water.  Add  4  or  5  drops  of  concentrated  hydrochloric  acid  and 
slowly  heat  to  boiling.    Set  aside  for  about  five  minutes,  then 
cool  and  neutralize  the  acid  in  the  solution  by  adding  a  concen- 
trated solution  of  sodium  carbonate  until  the  resulting  mix- 
ture is  just  alkaline  to  litmus  paper.    Test  this  with  Fehling's 
solution  according  to  above  directions.   Account  for  the  result. 

[97] 


136.  The  test  for  starch.    The  presence  of  starch  is  best 
shown  through  its  action  on  iodine  (§  104). 

137.  Alcohol.    The  formation  of  alcohol  and  carbon  dioxide 
from  glucose  may  be  shown  as  follows :  About  100  g.  of  glu- 
cose is  dissolved  in  a  liter    of  water  in  flask  A  (Fig.  52). 
This  flask  is  connected  with  the  bottle  B,  which  is  partly 
filled  with  limewater.    The  tube  C  contains  solid  sodium  hy- 
droxide.   A  little  baker's  yeast  is  now  added  to  the  solution 
in  flask  A,  and  the  apparatus  is  connected  as  shown  in  the 
figure.    If  the  tempera- 
ture   is   maintained    at 

about  30°,  the  reaction 
soon  begins,  but  it  will 
not  be  completed  for 
several  days.  The  ex- 
periment can  be  finished 
at  a  later  laboratory 
period.  The  bubbles  of 
gas  escape  through  the 
limewater  in  B.  A  pre-  FIG.  52 

cipitate  of  calcium  car- 
bonate soon  forms  in  the  limewater,  showing  the  presence 
of  carbon  dioxide.  The  sodium  hydroxide  in  tube  C  prevents 
the  carbon  dioxide  in  the  air  from  acting  on  the  limewater. 
The  alcohol  remains  in  the  flask  A.  Connect  the  flask  with 
a  condenser  (Fig.  24)  and  distill  about  10  cc.  of  the  solution. 
Pour  this  into  an  evaporating-dish  and  see  whether  it  will 
burn.  Test  a  few  drops  with  the  iodoform  test  (§  124). 

138.  Acetic  acid.    a.  Place  about  10  g.  of  sodium  acetate  in 
a  small  flask  (about  250  cc.)  provided  with  a  stopper  and 
delivery  tube.    Pour  5  cc.  of  concentrated  sulfuric  acid  into 
an  equal  volume  of  water,  and  add  the  solution  to  the  flask. 
Arrange  the  apparatus  so  that  the  delivery  tube  just  touches 
the  surface  of  a  little  water  contained  in  a  beaker,  and  then 
gently  heat  the  mixture  in  the  flask  until  it  has  boiled  for  a 
short  time  and  several  drops  of  distillate  have  passed  over  into 
the  water  in  the  beaker. 

[98] 


b.  Notice  the  odor  of  the  distillate.  Test  its  reaction  with 
litmus.  Add  sodium  hydroxide  to  the  distillate  until  it  is 
just  neutral,  and  then  pour  in  a  few  drops  of  a  solution  of 
ferric  chloride  (FeCl3).  The  solution  becomes  blood  red,  and 
this  reaction  constitutes  a  test  for  acetic  acid. 

139.  Esters.    Place  1  or  2  g.  of  sodium  acetate  in  a  test 
tube.    Add  1  or  2  cc.  of  alcohol  and  about  1  cc.  of  concen- 
trated sulfuric  acid,  and  warm  gently.    The  product  of  the 
reaction  is  ethyl  acetate  (CH3CO2C2H6).    Notice  the  pleasant 
odor.    The  formation  of  this  substance  with  its  characteristic 
odor  may  be  used  as  a  test  for  an  acetate  or  for  acetic  acid. 

140.  Determination  of  the   acidity  of   vinegar.    Introduce 
5  cc.  of  vinegar  into  a  small  beaker  and  dilute  with  about 
50  cc.  of  water.    Add  2  drops   of    an    alcoholic    solution  of 
phenolphthalein  and  then  slowly  run  in  a  solution  of  sodium 
hydroxide  until  the  solution  is  neutral,  as  in  §  71. 

If  a  solution  of  sodium  hydroxide  of  known  strength  is 
used,  it  is  evident  that  one  may  in  this  way  determine  the 
percentage  of  acetic  acid  in  the  vinegar.  The  laws  of  most  of 
the  states  require  a  minimum  of  4  per  cent  of  acetic  acid  in 
all  vinegars. 

141.  Saponification  and  the  preparation  of  soap.  Add  10  cc. 
of  alcohol  to  4  or  5  g.  of  cottonseed  oil  in  an  evaporating- 
dish.    To  the  resulting  mixture  add  about  1  cc.  of  a  50  per 
cent  solution  of  sodium  hydroxide  in  water.    Evaporate  care- 
fully, stirring  the  mixture  constantly  until  the  odor  of  alcohol 
can  no  longer  be  detected.   Write  the  equation  for  the  reaction 
on  the  supposition  that  the  oil  is  composed  of  palmitin.    What 
remains  in  the  dish  ?     Add  75  cc.  of  cold  water,  stir  well, 
and  filter.    Add  a  little  of  the  resulting  filtrate  to  test  tubes 
containing  solutions  of  calcium  sulfate,  magnesium  sulfate,  and 
sodium  sulfate  respectively.     Explain  the  results.    Account 
for  the  fact  that  soaps  do  not  lather  freely  in  hard  waters. 


[99] 


CHAPTER  XXIV 
THE  PHOSPHORUS  FAMILY 

142.  Phosphorous  acid.    Pour  into  a  test  tube  about  0.5  cc. 
of  phosphorus  trichloride   (hood)   and  add  a  little  water,  a 
drop  at  a  time.    What  gas   is   evolved  (R)  ?    Finally,  add 
about  5  cc.   of  water,  pour  the  liquid  into  an  evaporating- 
dish,  and  evaporate  to  a  sirupy  mass.    Dilute  this  with  a 
little  water,  transfer  the  solution  to  a  test  tube,  and  add  a 
few  drops  of  a  solution  of  silver  nitrate.    Boil  the  resulting 
mixture.    A  black  precipitate  forms.    Explain. 

143.  Phosphoric  acid.    a.  Add,  in  very  small  portions,  about 
1  g.  of  phosphorus  pentoxide  to  15  cc.  of  cold  water.    To  1  cc. 
of  the  solution  add  a  solution  of  silver  nitrate  until  a  precipi- 
tate forms.    The  precipitate  is  AgPO3.    Explain.    Pour  the 
remainder  of  the  original  solution  into  an  evaporating-dish 
and  evaporate  to  a  sirupy  mass.    Add  about  lOcc.  of  water 
and  carefully  neutralize  the  resulting  solution  with  sodium 
carbonate.    Na2HPO4  is  formed.    Test  the  solution  with  silver 
nitrate.    The  equation  for  the  reaction  is  as  follows : 

Na2HPO4  +  3  AgNO3  =  Ag3PO4  +  2  NaNO3  +  HNO8 

b.  Test  a  solution  of  normal  sodium  phosphate  with  litmus 
paper.    Explain.    Add  a  little  silver  nitrate  (R). 

c.  Add  a  few  drops  of  magnesia  mixture  to  a  solution  of 
disodium  phosphate.    (Magnesia  mixture  is  prepared  by  dis- 
solving 1  g.  of  magnesium  chloride  crystals  and  3  g.  of  am- 
monium   chloride  in   10  cc.   of   water  and    adding   1  cc.   of 
ammonium  hydroxide  solution.)   The  precipitate  has  the  com- 
position MgNH4PO4.    Advantage  is  taken  of  this  reaction  in 
determining  the  amount  of  magnesium  in  various  substances. 

[100] 


d.  With  litmus  paper  test  the  reaction  of  disodium  phos- 
phate and  of  silver  nitrate.    Now  add  a  solution  of  silver 
nitrate  to  one  of  disodium  phosphate  until  an  appreciable 
precipitate   has  been  formed,   and   test   the    mother   liquor. 
How  do  you  account  for  the  reaction  ? 

e.  Add  a  few  drops  of  a  solution  of  disodium  phosphate  to 
5  cc.  of  a  solution  of  ammonium  molybdate  (side  shelf).    A 
bright  yellow  precipitate  of  complex  formula  slowly  forms 
(§  88,  6). 

144.  Pyrophosphates.    Apply  a  gentle  heat  to  1  or  2  g.  of 
disodium  phosphate  placed  in  a  porcelain  crucible.    Gradually 
increase  the  heat  to  the  full  extent  and  continue  the  heating 
for  about  ten  minutes.    When  the  crucible  has  cooled,  dissolve 
the  product  in  cold  water  and  test  the  solution  with  silver 
nitrate  solution.    Compare  with  the  product  obtained  by  add- 
ing silver  nitrate  to  disodium  phosphate  which  has  not  been 
heated.  Explain. 

145.  Metaphosphates.    a.  Repeat  the  experiment  of  §  144, 
using  microcosmic  salt  (NaNH4HPO4)  instead  of  disodium 
phosphate.     (This  salt,  on  being  heated,   acts  just  as  does 
NaH2PO4.)    Explain. 

b.  Compare  the  results  obtained  by  adding  silver  nitrate 
to  solutions  of  the  salts  of  each  of  the  three  phosphoric  acids. 

c.  Make  a  loop  on  the  end  of  your  platinum  wire  and  in 
this  form  a  bead  like  the  borax  bead  (§  151,  a),  using  micro- 
cosmic  salt  instead  of  borax  (R).   Of  what  will  the  clear  bead 
consist  ?    What  color  is  given  to  the  bead  by  rubbing  it  in  a 
little  copper  oxide  and  reheating  (R)  ?    Test  the  color  occa- 
sioned by  a  compound  of  cobalt  and  one  of  chromium.    What 
is  formed  when  sodium  metaphosphate  dissolves  these  oxides  ? 
Will  sodium  nitrate  form  similar  compounds?   Why?   Could 
sodium  metaphosphate  be  used  as  a  flux  in  brazing  ? 

146.  Arsenic,    a.  Note  the  physical  properties  of  arsenic. 
Heat  (hood)  a  bit  of  it  half  as  large  as  a  gram  of  wheat,  on 
charcoal,  with  a  blowpipe.    Explain. 

b.  Seal  the  end  of  a  piece  of  glass  tubing  about  10  cm.  in 
length  and  6  or  7  mm.  in  diameter.     Introduce  into  it  an 

[101] 


amount  of  arsenious  oxide  equal  in  bulk  to  a  grain  of  wheat. 
Cover  this  to  a  depth  of  2  or  3  cm.  with  somewhat  finely 
powdered  charcoal  which  has  been  strongly  heated  in  a  por- 
celain crucible.  See  that  the  inner  surface  of  the  tube  above 
the  charcoal  is  perfectly  clean.  Incline  the  tube  and  heat  the 
upper  portion  of  the  charcoal  to  a  high  temperature ;  then, 
while  maintaining  the  charcoal  at  this  temperature,  gradually 
bring  the  lower  part  of  the  tube  also  into  the  flame  (Fig.  53). 
The  upper  part  of  the  tube  must  be  kept  as  cool  as  possible. 
The  arsenious  oxide  is  changed  into  a  vapor,  which  passes  over 
the  hot  charcoal.  Account  for  the  result  (R).  Cut  the  tube  as 
near  the  bottom  as  pos- 
sible and  remove  the 
charcoal;  then,  inclin- 
ing the  tube,  apply  a 
very  gentle  heat  to  the 
portion  of  it  which  con- 
tains the  dark  mirror. 
Note  that  small  white 
crystals  are  slowly  de- 
posited in  the  colder 
portions  of  the  tube 
(R).  Note  their  form. 

c.  Marsh's  test  (hood).  The  arsine  formed  in  this  experi- 
ment is  very  poisonous,  and  great  care  must  be  taken  to 
prevent  its  escape  into  the  air  of  the  laboratory. 

Arrange  an  apparatus  according  to  Fig.  54,  consisting  of 
the  generator  A,  a  calcium  chloride  drying-tube  B,  and  a  clean 
hard-glass  tube  C  about  30  cm.  long  and  8  mm.  in  diameter, 
drawn  out  to  a  jet  at  the  end.  (Use  the  blast  lamp  in  making 
the  jet.)  Generate  hydrogen  by  the  usual  method  and,  after 
taking  the  general  precautions  (§  24,  d),  ignite  it  as  it  escapes 
from  the  glass  jet  D.  Sufficient  acid  is  added  from  time  to 
time  to  cause  a  gentle  evolution  of  the  gas.  Now  apply  a 
strong  heat  to  the  hard-glass  tube  at  a  place  near  its  middle, 
using  the  wing-top  burner.  After  a  few  minutes  note  whether 
any  deposit  forms  just  beyond  the  heated  portion  of  the  tube. 

[102] 


FIG.  53 


If  none  forms,  the  materials  are  free  from  arsenic.  Now  ask 
the  instructor  to  add  to  the  generator  2  drops  of  a  dilute 
solution  of  arsenious  oxide  in  hydrochloric  acid,  rinsing  it 
down  the  funnel  tube  with  a  little  water.  Continue  the  heat- 
ing of  the  hard-glass  tube  at  the  same  place.  Note  the  deposit 
formed  on  the  sides  of  the  tube.  Withdraw  the  heat  and  hold 
the  lid  of  a  porcelain  crucible  in  the  flame.  A  black  deposit 
of  arsenic  forms.  Test  the  solubility  of  this  in  a  solution  of 
sodium  hypochlorite.  Note  the  results.  Cut  the  tube  contain- 
ing the  deposit  so  as  just  to  remove  the  jet,  and,  inclining  it, 


O 


B 


FIG.  54 

apply  a  gentle  heat,  as  in  b.  Account  for  the  results  and  write 
the  equations  for  all  the  reactions  involved  in  the  experiment. 
147.  Arsenious  oxide.  Place  a  very  little  arsenious  oxide  in 
a  test  tube  and  add  a  few  cubic  centimeters  of  a  solution  of 
sodium  hydroxide,  heating  the  solution  gently.  Does  the 
oxide  dissolve  (R)  ?  What  do  you  infer  as  to  the  character 
of  arsenious  oxide  ?  Save  the  solution  for  §  148.  Repeat  the 
experiment,  substituting  hydrochloric  acid  for  the  sodium 
hydroxide  (R).  What  would  this  experiment  suggest  as  to 
the  nature  of  arsenious  oxide  ?  Set  the  test  tube  aside  and 
note  the  crystals  which  separate.  They  are  As0O3.  What  light 
does  this  throw  on  the  nature  of  arsenious  oxide  ? 

[103] 


148.  Arsenious  sulfide.    Pass  a  current  of  hydrogen  sulfide 
through  the  two  solutions  just  described.    What  difference  in 
conduct  do  you  note  ?    What  does  the  yellow  color  indicate  ? 
Is  arsenious  sulfide  soluble  in  water  ?    Why  does  it  not  pre- 
cipitate in  the  first  solution  ?    Now  add  hydrochloric  acid  to 
this  solution  until  it  is  acid  in  reaction.    What  change  takes 
place  ?    How  do  you  explain  it  ?    What  is  the  precipitate  ? 
Collect  it  by  filtering  through  a  small  filter,  and  transfer  the 
precipitate  to  a  test  tube.    Cover  it  with  concentrated  hydro 
chloric  acid,  add  a  bit  of  potassium  chlorate,  and  heat  to 
boiling  (hood).    The   sulfide   is  converted  into  the  soluble 
arsenic  acid,  in  accordance  with  the  following  equation: 

As2S3  + 10  Cl  +  8  H20  =  2  H3AsO4  +  3  S  + 10  HC1 

What  is  the  source  of  the  chlorine  ?  What  is  the  function 
of  the  potassium  chlorate  ?  Now  heat  the  solution  containing 
the  arsenic  acid  until  the  chlorine  is  all  expelled,  add  ammo- 
nium hydroxide  until  the  solution  is  alkaline,  and  then 
enough  more  to  increase  the  volume  by  one  third.  Filter 
unless  the  liquid  is  clear.  Add  a  few  drops  of  magnesia  mix- 
ture and  shake  the  solution  vigorously.  The  precipitate  is 
MgNH4AsO4.  Compare  with  the  action  of  magnesia  mixture 
on  phosphoric  acid  (§  143,  c). 

Tabulate  the  names  and  the  formulas  for  the  acids  of  ar- 
senic. How  do  they  compare  in  composition  with  those  of 
phosphorus  ? 

149.  Antimony,    a.  Repeat  §  146,  a,  substituting  antimony 
for  the  arsenic.    Note  the  results. 

b.  Repeat  §  146,  c,  substituting  2  or  3  drops  of  a  solution 
of  a  compound  of  antimony  for  the  arsenious  oxide.   Compare 
the  results  of  the  two  experiments  (R). 

c.  Introduce  a  bit  of  antimony  no  larger  than  a  grain  of 
wheat  into  a  test  tube,  and  add  about  2  cc.  of  hydrochloric 
acid.    Does    the    antimony    dissolve  ?    Why  ?    Add    2    or   3 
drops  of  nitric  acid  (hood).    After  the  antimony  has  been 
dissolved,  pour  the  solution  into  a  beaker  containing  about 
100  cc.  of  water  (R).    Half  fill  a  test  tube  with  the  resulting 

[104] 


mixture,  add  hydrochloric  acid,  a  drop  at  a  time,  until  the 
solution  just  clears  (R).  Now  pass  hydrogen  sulfide  through 
the  clear  solution  until  an  orange-colored  solid  forms  (R). 

d.  Boil  a  small  amount  of  powdered  antimony  with  con- 
centrated nitric  acid  (hood)  until  it  is  changed  into  a  white 
powder.  Dilute  the  mixture  with  water,  filter,  and  wash  the 
residue  on  the  filter  paper  with  water.  Convince  yourself 
that  the  residue  is  not  a  nitrate.  It  is  a  hydrated  oxide  of 
variable  formula,  but  may  be  considered  to  have  the  formula 
Sb2O3.  Dissolve  a  portion  of  it  in  sodium  hydroxide  (R). 
Dissolve  another  portion  in  hydrochloric  acid  (R). 

150.  Bismuth,  a.  Repeat  §  146,  a,  substituting  bismuth 
for  arsenic.  Contrast  the  action  of  heat  on  arsenic,  on  anti- 
mony, and  on  bismuth. 

b.  Repeat  §  149,  c,  substituting  bismuth  for  antimony  and 
using  nitric  acid  alone  as  the  solvent. 


[105] 


CHAPTER  XXV 

• 

BORON  AND  SILICON 

151.  Borax  bead.    a.  Make  a  little  loop  on  the  end  of  a 
platinum  wire  and  heat  it  to  redness  in  a  Bunsen  flame  ;  then 
quickly  bring  the  loop  in  contact  with  some  borax  and  reheat. 
The  borax  adhering  to  the  loop  will  swell  up  (why?)  and 
finally  form  a  clear  glass  bead.    Now  dip  the  bead  into  a 
solution  of  a  cobalt  compound  and  reheat  thoroughly  (R). 

b.  Repeat   the    experiment,    substituting    a    compound    of 
chromium  for  the   cobalt   solution. 

152.  Boric  acid.    a.  Dissolve  5  g.  of  borax  in  15  cc.  of  boil- 
ing water.    Test  the  solution  with  litmus  paper.    Explain. 
Add  to  the  hot  solution  5  cc.  of  concentrated  hydrochloric 
acid.  Cool  the  solution,  filter  off  the  precipitate  (R).  Compare 
the  precipitate  with  borax  as  to  solubility  in  alcohol. 

b.  Place  1  cc.  of  a  solution  of  borax  in  an  evaporating-dish 
and  add  a  few  drops  of  sulfuric  acid  (R)  and  2  or  3  cc.  of 
alcohol.    Set  fire  to  the  alcohol  and  watch  for  a  green  color  in 
the  flame.   This  is  a  test  for  boric  acid.    Will  borax  act  in  the 
same  way  if  no  sulfuric  acid  is  added  ? 

c.  Heat  a  little  boric  acid  in  an  iron  crucible  until  a  clear 
liquid  is  formed  (R). 

153.  Borax.    Prepare  a  concentrated  solution  of  borax  and 
add  a  few  drops  of  a  solution  of  silver  nitrate.    The  precipitate 
is  silver  borate  (R).    Add  a  few  drops  of  the  concentrated 
solution  of  borax  to  a  test  tube  half  full  of  water,  and  then 
test  with  silver  nitrate.     The  precipitate  is  silver  oxide  (R). 
Compare  it  with  the  precipitate  formed  by  the  action  of  very 
dilute  sodium  hydroxide  on  silver  nitrate  (R).    How  do  you 
account  for  the  different  actions  of  borax  on  silver  nitrate  ? 

[106] 


154.  Silica,  a.  Mix  about  half  a  gram  of  fine  sand  with 
3  or  4  times  its  weight  of  solid  sodium  hydroxide.  Place 
the  mixture  in  an  iron  dish,  and  heat  until  fusion  has  taken 
place  and  the  fused  mass  has  again  become  solid  (R).  Then 
dissolve  the  product  in  hot  water. 

&.  Pour  half  of  the  solution  into  a  test  tube  and  add  an 
excess  of  hydrochloric  acid.  Allow  it  to  become  perfectly 
cold.  What  is  the  jellylike  substance  ? 

c.  Transfer  the  remainder  of  the  solution  to  an  evaporating- 
dish,  acidify  with  hydrochloric  acid  (R),  evaporate  to  dryness, 
and  heat  the  dish  gently  with  the  bare  flame  (R).  Is  the 
product  soluble  in  water  ?  in  acids  ?  in  alkalis  ?  in  fused 
sodium  hydroxide  ? 


[107] 


CHAPTER  XXVI 
COLLOIDS 

155.  Preparation  of  colloids,     a.  Prepare  4  or  5  cc.  of   a 
saturated  solution  of  sulfur  in  alcohol  and  pour  it  into  100  cc. 
of  water.    The  sulfur  forms  a  fine  white  dispersion  that  will 
not  settle.    Filter  the  mixture  (?). 

b.  To  10  cc.  of  water  in  a  test  tube  add  3  drops  of  a  dilute 
solution  of  gold  chloride  (6  g.  crystallized  HAuCl4  •  3H2O  per 
liter).    Heat  to  boiling  and  add  1  or  2  cc.  of  a  dilute  solution 
of  formalin  as  a  reducing  agent  (3  cc.  commercial  formalin 
per  liter  of  water).    The  solution  assumes  a  purplish  color 
due  to  metallic  gold.    Compare  the  appearance  viewed   by 
transmitted  light  with  that  by  reflected  light. 

c.  To  100  cc.  of  water  add  2  or  3  drops  of  the  solution  of 
gold  chloride.    After  thorough  shaking,  add  a  few  drops  of  a 
dilute  solution  of  tannin  (about  0.1  g.  per  liter)  and  heat  to 
near  boiling.    If  the  color  is  too  faint  repeat  the  addition  of 
the  two  reagents.     How  does  the  color  compare  with  that 
obtained  in  b  ?    How  do  you  account  for  the  difference  ?    The 
tannin  acts  both  as  a  reducing  agent  and  a  protecting  colloid. 
To  the  red  solution  add  1  cc.  of  a  concentrated  solution  of 
sodium  chloride.    What  change  do  you  notice  ? 

156.  Preparation  of  colloids  by  double  decomposition,  a.  Pre- 
pare a  little  antimony  sulfide  by  dissolving  a  small  grain  of 
antimony  oxide  or  antimony  chloride  in  5  cc.  of  dilute  hydro- 
chloric acid  and  adding  a  solution  of  hydrogen  sulfide.    The 
orange-colored  precipitate  has  the  composition  Sb2Sg.    When 
the  precipitate  has  settled,  is  the  liquid  colored  ?    Set  the  tube 
containing  it  aside  for  reference. 

[108] 


b.  Dissolve  a  little  tartar  emetic  (which  contains  antimony 
as  a  salt  of  a  weak  acid),  not  exceeding  a  wheat  grain  in  bulk, 
in  about  100  cc.  of  water,  and  add  a  few  drops  of  a  solution 
of  hydrogen  sulfide,  taking  care  to  avoid  an  excess.  (The 
resulting  solution  should  not  have  the  odor  of  hydrogen  sul- 
fide.) What  change  do  you  notice?  Filter  a  little  of  the 
solution.  Is  the  filtrate  colorless  ?  Is  a  precipitate  left  on 
the  filter  ?  Boil  2  or  3  cc.  of  the  filtrate  and  again  filter.  Does 
boiling  cause  the  antimony  sulfide  to  separate  ?  Reserve  the 
remainder  of  the  solution  for  §§  158,  159,  160. 

157.  Preparation  of  colloids  by  hydrolysis,    a.  To  100  cc. 
of  water  in  a  beaker  add  enough  of  a  solution  of  ferric  chlo- 
ride to  give  a  decidedly  yellow  color.    Slowly  heat  the  solution 
to  boiling  and  note  the  change  in  color.    The  salt  is  hydro- 
lyzed  with  the  formation  of  colloidal  ferric  oxide.    View  the 
color  by  reflected  and  then  by   transmitted  light.    Compare 
the  latter  with  the  color  of  ferric  oxide  precipitated  from  dilute 
ferric  chloride  by  ammonium  hydroxide.    Save  the  solution 
for  §§  157,  158,  159. 

b.  Dissolve  2  or  3  g.  of  sodium  acetate  in  50  cc.  of  water. 
Now  add  1  cc.  of  a  solution  of  ferric  chloride.  What  new  salts 
should  be  formed  ?  Should  either  of  these  undergo  extensive 
hydrolysis  ?  Do  you  obtain  a  precipitate  ?  What  is  the  color 
of  the  solution  ?  This  is  largely  due  to  the  presence  of  col- 
loidal ferric  oxide.  Save  this  solution  for  §§  158,  159,  160. 

158.  Coagulation  of  colloids.    Heat  to  boiling  5  cc.  of  the 
colloidal  solutions  obtained  in  156,  &,  and  157,  a  and  b.    Is 
the  colloid  precipitated  ?    To  5  cc.  of  each  add  about  half  a 
gram  of  ammonium  chloride  or  calcium  chloride  and  shake 
the  solution  vigorously.    Do  any  of  the  solutions  give  a  pre- 
cipitate ?    To  5  cc.  of  each  add  1  drop  of  concentrated  sul- 
furic  acid.    Would  you  expect  this  reagent  to  precipitate 
ferric  oxide  ?   Explain.   Remembering  that  fermentation  often 
produces  acids,  how  can  you  explain  the  curdling  of  milk  on 
souring  ? 

159.  Protective  colloids.   Dissolve  a  few  small  pieces  of  agar 
(a  colloidal  material  resembling  gelatin)  in  about  10  cc.  of 

[109] 


hot  water.  Add  equal  portions  of  this  solution  to  portions 
of  the  colloidal  solutions  of  antimony  sulfide  and  of  ferric 
oxide.  With  these  solutions  repeat  the  experiments  of  §  158. 
Does  the  presence  of  the  agar  interfere  with  coagulation  ? 
It  is  often  the  case  that  the  presence  of  organic  colloids  inter- 
feres with  precipitation  tests  in  chemical  analysis.  Such 
colloids  are  then  called  protective  colloids. 

160.  Precipitation  of  one  colloid  by  another.   Sometimes  one 
colloid  will  act  upon  another,  causing  mutual  precipitation. 
Treat  some  of  the  colloidal  solution  of  iron  oxide  with  an 
equal  volume  of  the  antimony  sulfide,  taking  care  that  no 
free  hydrogen  sulfide  is  present.    Does  a  precipitate  form  ? 
Filter  some  of  the  solution.    Is  the  filtrate  red  like  the  iron, 
or  orange  like  the  antimony  ? 

161.  Adsorption,   a.  Fold  a  small  filter  paper  to  fit  a  funnel 
and  place  upon  it  a  layer  of  bone  black  2  or  3  cm.  thick.   Wet 
it  thoroughly  with  water  and  let  it  drain.    Prepare  50  cc.  of 
a  solution  decidedly  colored  by  blue  litmus,  congo  red,  or 
some  other  dye.    Slowly  filter  the  colored  solution  through 
the  bone  black,  repeating  several  times.   Is  the  color  removed  ? 

b.  Repeat  the  experiment,  using  a  dilute  solution  of  am- 
monia instead  of  the  dye.  Can  you  detect  ammonia  in  the 
filtrate  ?  How  does  it  affect  red  litmus  ? 

162.  Emulsions,    a.  In  a  250-cc.  flask  pour  5  cc.  of  a  5  per 
cent  solution  of  hard  soap  in  water.    Add  benzene  in  suc- 
cessive portions  as  follows:   5,  10,  20,  20,  20,  20,  30,  50  cc., 
shaking  the  stoppered  flask  vigorously  after  each  addition. 
A  stiff  emulsion  should  be  obtained  consisting  of  about  97  per 
cent  benzene  emulsified  in  3  per  cent  of  water  by  a  very 
little  soap. 

b.  Shake  together  20  cc.  of  raw  linseed  oil  and  20  cc.  of 
water,  adding  the  water  in  small  portions  at  a  time.  The  oil 
always  contains  enough  rosin  to  act  as  the  emulsifying  agent. 

163.  Gels.    a.  Break  up  some  dry  gelatin  or  agar  into  very 
small  pieces  and  place  them  in  a  test  tube  to  a  depth  of  about 
2  cm.   Add  10  cc.  of  water  and  heat  nearly  to  boiling  until  a 
clear  solution  is  obtained.    Immerse  the  test  tubes  in  ice  water 

[110] 


for  a  time.  What  change  do  you  note  ?  Now  put  the  tube 
in  hot  water.  Does  the  gel  liquefy  ?  Once  more  cool.  Does 
the  gel  form  again  ?  What  is  such  a  gel  called  ? 

b.  To  10  cc.  of  water  glass  of  density  1.1  add  10  drops  of 
concentrated   sulfuric   acid.     Shake  the   solution   thoroughly 
and  set  aside  if  necessary  until  the  next  period.    Tap  the 
test  tube  gently  and  note  the  vibrations  of  the  firm  gel  of 
silicic  acid. 

c.  Immerse  the  tube  containing  the  gel  in  hot  water.   Does 
the  gel  liquefy  ?    With  a  glass  rod  dig  out  a  little  of  the 
gel  and  see  if  you  can  dissolve  it  in  water.    What  is  such  a 
gel  called  ? 


CHAPTER  XXVII 
GENERAL  METHODS  FOR  PREPARATION  OF  COMPOUNDS 

164.  General  methods,  a.  Chlorides.  How  can  you  prepare 
zinc  chloride  from  zinc  ?  from  zinc  carbonate  ?  from  zinc 
oxide  (R)  ?  How  can  you  prepare  calcium  chloride  from 
calcium  carbonate  ?  What  chlorides  are  insoluble  ?  Prepare 
small  amounts  of  each  (R)  and  note  their  physical  properties. 

b.  Sulfides.    Describe  the  different  ways  in  which  sulfides 
have  been  prepared  in  the  laboratory.    Devise  a  method  for 
preparing  the  following  insoluble  sulfides :  lead  sulfide,  silver 
sulfide,   antimony    sulfide,    zinc    sulfide,    manganese    sulfide. 
Prepare  (hood)  small  amounts  of  each  in  test  tubes  (R). 

c.  Nitrates.   How  can  you  prepare  copper  nitrate  from  cop- 
per ?  from  copper  carbonate  ?  from  copper  hydroxide  (R)  ? 
Are  any  of  the  nitrates  insoluble  ?   What  is  the  effect  of  heat 
on  copper  nitrate  ?  on  lead  nitrate  ?  on  ammonium  nitrate  ? 

d.  Nitrites.    Recall  the  preparation  of  sodium  nitrite  (R). 

e.  Sulfates.    The  following  sulfates  have  been  prepared 
in  previous  exercises:    zinc  sulfate,  sodium  sulfate,  copper 
sulfate,  iron  sulfate.    Write  the  equations  for  the  reactions 
involved  in  the  preparation  of  each.    What  sulfates  are  insol- 
uble ?    Prepare  small  amounts  of  each,  noting  the  color  and 
writing  the  equations  for  the  reactions  involved  in  each  case. 

/.  Sulfites.  Recall  the  method  used  in  the  preparation  of 
sodium  sulfite.  How  can  you  distinguish  between  sulfates 
and  sulfites? 

g.  Carbonates.  What  carbonates  are  soluble  ?  Give  a  gen- 
eral method  for  the  preparation  of  the  soluble  carbonates  ; 
give  a  general  method  for  the  preparation  of  the  insoluble 
carbonates.  Prepare  calcium  carbonate  by  two  different 

[112] 


methods  (R).  How  could  you  prepare  from  calcium  car- 
bonate the  following  compounds:  calcium  chloride,  calcium 
sulfate,  calcium  nitrate,  calcium  oxide  (R)  ?  Why  are 
carbonates  so  readily  decomposed  by  acids  ? 

h.  Phosphates.  Recall  the  formulas  and  the  names  of  the 
three  phosphoric  acids,  also  the  action  of  silver  nitrate  on 
their  salts. 

/.  Oxides.  Burn  a  bit  of  magnesium  wire  1  or  2  cm.  in 
length  (R).  Heat  a  crystal  of  lead  nitrate  in  a  test  tube  (R). 
Heat  a  small  piece  of  limestone  on  a  wire  gauze  (R). 

j.  Hydroxides.  Hold  a  piece  of  lime  the  size  of  a  walnut  in 
water  for  a  few  seconds,  then  place  it  on  a  watch  glass  and 
set  it  aside  for  a  half  hour.  Note  the  change  (R).  Pour  into 
separate  test  tubes  about  1  cc.  of  a  solution  of  each  of  the 
following  compounds :  ferric  chloride  (FeCl3) ;  magnesium 
sulfate  (MgSO4) ;  copper  sulfate  (CuSO4).  Add  to  each 
solution  2  drops  of  a  solution  of  sodium  hydroxide  (R). 
Now  add  a  few  drops  of  hydrochloric  acid  to  each  test 
tube,  and  explain  the  results. 

165.  Equilibrium,  a.  To  3  or  4  cc.  of  a  solution  of  barium 
chloride  add  dilute  sulfuric  acid,  drop  by  drop,  as  long  as  a 
precipitate  is  produced.  Is  the  reaction  a  complete  one  ?  In 
a  similar  way,  to  a  solution  of  calcium  chloride  add  a  solu- 
tion of  oxalic  acid  as  long  as  calcium  oxalate  (CaC2O4)  is 
precipitated.  To  test  the  completeness  of  the  reaction,  filter 
off  the  precipitate  and  add  ammonium  hydroxide  to  the  fil- 
trate until  it  is  alkaline  in  reaction.  Does  an  additional  pre- 
cipitate form  ?  Can  it  be  ammonium  oxalate  ?  What  is  it  ? 
Why  did  it  not  appear  before  the  ammonium  hydroxide  was 
added  ?  If  ammonium  oxalate  is  added  in  excess  to  calcium 
chloride,  is  the  reaction  complete  ?  Why  is  it  incomplete 
when  the  free  acid  is  used  ? 

b.  Make  a  list  of  about  a  dozen  acids  which  you  have 
studied.  What  insoluble  salts  do  they  form  ?  Prepare  an  in- 
soluble salt  of  each  (if  it  forms  one)  and  try  its  solubility  in 
hydrochloric  acid.  Which  ones  should  be  insoluble  ?  Why  ? 
Do  your  experiments  confirm  your  predictions? 

[113] 


CHAPTER  XXVIII 
THE  ALKALI  METALS 

NOTE.  Some  preliminary  experiments  with  sodium  and  sodium 
hydroxide  have  been  given  in  Chapter  XI.  The  student  should  review 
these  carefully  in  connection  with  the  present  chapter. 

166.  Action  of  sodium  hydroxide  upon  soluble  salts  of  other 
metals.    Try  the  effect  of  the  solution  of  sodium  hydroxide 
upon  a  soluble  salt  of  each  of  the  following:  barium,  potas- 
sium, magnesium,  zinc,  copper  (R).    In  each  case  add  a  few 
drops  at  first,  and  if  a  precipitate  forms,  add  an  excess  to  see 
whether  the  precipitate  redissolves.    Make  a  list  of  the  hy- 
droxides soluble  in  water,  and  another  list  of  those  which 
redissolve  in  an  excess  of  sodium  hydroxide. 

167.  Preparation  of  sodium  chloride.   Dissolve  5  g.  of  sodium 
carbonate  in  20  cc.  of  water.    Prepare  common  salt  from  this 
(R).    How  can  you  be  sure  that  the  product  contains  no  un- 
changed sodium  carbonate  ?    Describe  the  method  (R).    Treat 
some  of  the  salt  so  prepared  with  sulfuric  acid.    What  gas  is 
evolved  (R)  ? 

168.  Purification   of   common   salt.    The   impurities   most 
likely  to  be  present  are  sulfates  and  chlorides  of  calcium  and 
of  magnesium.    Dissolve  about  25  g.  of  common  salt  in  100 
or  150  cc.  of  water  and  heat  to  boiling.    Add  a  solution  of 
barium  hydroxide,  1  cc.  at  a  time,  as  long  as  you  are  sure 
that  the  precipitate  is  increasing  (R).    Without  filtering,  add 
a  concentrated  solution  of  sodium  carbonate,  drop  by  drop,  as 
long  as  it  produces  a  precipitate.    If  the  solution  becomes  too 
turbid  for  you  to  be  sure  of  this,  filter  it.    What  is  precipi- 
tated (R)  ?    When  the  precipitation  is  complete,  filter.    What 
impurity  may  now  be  present  ?    Add  dilute  hydrochloric  acid 

[114] 


until  the  solution  is  distinctly  acid  (R),  and  then  evaporate 
to  dryness  in  an  evaporating-dish,  using  a  very  small  flame 
toward  the  end  of  the  process.  Test  the  product  for  each  of 
the  original  impurities.  Expose  some  of  it  to  the  air  for  a 
number  of  days.  Does  it  become  moist? 

169.  Sodium  thiosulfate.    a.  Dissolve  12  g.  of  crystallized 
sodium  sulfite  (Na2SO3  •  7  H2O)  in  25  cc.  of  water  and  place 
the  solution  in  a  small  flask  provided  with  a  stopper  through 
which  a  short  piece  of  glass  tubing  has  been  inserted.    Add 
1.5  g.  of  finely  powdered  roll  sulfur  (not  flowers  of  sulfur) 
and  gently  boil  the  mixture  until  the  sulfur  has  nearly  all 
dissolved.    This  may  require  an  hour  or  more.    Why  do  the 
stopper  and  the  tube  largely  prevent  oxidation  by  the  air? 
Will  the  reaction  be  more  rapid  if  the  solution  is  boiled  vig- 
orously rather  than  very  gently  ?    Why  ?    Filter  the  solution 
and  set  it  aside  to  crystallize.    If  no  crystals  form  by  the 
next  day,  add  a  very  small  crystal  of  the  thiosulfate  to  start 
crystallization.    Explain.    Filter  off  the  crystals  and  dry  them 
on  filter  paper.    Are  they  efflorescent  or  deliquescent  ? 

b.  To  2  or  3  cc.  of  the  mother  liquor  add  a  little  hydro- 
chloric acid  (R).  Prepare  a  little  silver  chloride  by  pre- 
cipitation and,  after  washing  it  with  water,  add  to  it  1  or 
2  cc.  of  the  thiosulfate  solution.  Does  it  dissolve  ?  What 
use  is  made  of  this  reaction  ?  Test  the  solubility  of  one  or 
two  crystals  of  iodine  in  water  and  then  add  a  little  thio- 
sulfate solution  (R).  How  is  this  reaction  used  in  chemical 
analysis  ? 

170.  Potassium   nitrate.    Give   the   sources   of   potassium 
nitrate.    Put  into  a  small  beaker  the  theoretical  amounts  of 
sodium  nitrate  and  potassium  chloride  necessary  to  prepare 
10  g.  of  potassium  nitrate.    What  are  these  weights  ?    Add 
20  cc.  of  water  and  boil  gently  (R).    The  sodium  chloride 
formed,  being  little  more  soluble  in  hot  water  than  in  cold,  soon 
begins  to  separate.    Stir  the  solution  and  continue  the  heat- 
ing until  the  volume  of  the  liquid  is  reduced  about  one  half. 
Quickly  filter  the  hot  solution  and  set  the  filtrate  aside  until 
cold.    Dry  the  crystals  deposited  in  the  filtrate  by  pressing 

[115] 


them  between  filter  papers.  Finally,  recrystallize  them  from 
as  little  hot  water  as  possible.  Prove  the  composition  of  the 
purified  crystals.  If  the  solution  is  allowed  to  cool  before 
filtering,  will  the  preparation  be  a  success  ?  Why  ? 

171.  Potassium  bromate  and  potassium  bromide.    Dissolve 
10  g.  of  potassium  hydroxide  in  15  cc.  of  water  in  a  100-cc. 
flask.    Weigh  out  (in  the  hood)  the  amount  of  bromine  nec- 
essary to  convert  this  into  a  mixture  of  the  bromide  and 
bromate  (R).    If  it  is  remembered  that  the  density  of  liquid 
bromine  is  almost  exactly  3,  the  liquid  may  be  measured 
instead  of  weighed.    Add  the  bromine,  a  drop  at  a  time,  to 
the  hydroxide  solution,  shaking  the  flask  after  the  addition 
of  each  drop.    If  the  solution  becomes  very  warm,  cool  it  by 
immersing  the  flask  in  cold  water.   If  the  hydroxide  is  impure, 
there  will  be  an  excess  of  bromine,  which  will  color  the  solu- 
tion.   In  this  case  add  an  excess  of  the  hydroxide,  a  drop  at 
a  time,  until  the  solution  just  becomes  colorless.    Cork  the 
flask  and  set  it  aside  until  the  next  laboratory  period,  then 
filter  off  the  bromate.    Evaporate  the  filtrate  to  dryness  in  an 
evaporating-dish  and  heat  to  a  high  temperature.    What  is 
the  residue  ?    Dry  the  bromate  and  heat  a  portion  of  it  in  a 
hard-glass  test  tube.    What  gas  is  evolved  (R)  ? 

172.  Potassium  chlorate  and  potassium  chloride.    Prepare 
potassium  chlorate  and  potassium  chloride  by  a  method  analo- 
gous to  that  of  §  171,  generating  chlorine  as  in  §  94,  e,  and 
passing  it  into  the  solution  of  potassium  hydroxide  till  the 
liquid  no  longer  feels  soapy.    See  also  §  107,  a. 

173.  Hydrolysis  of  salts.    Test  with  litmus  paper  solutions 
of  the  following  compounds :  sodium  carbonate,  sodium  hydro- 
gen carbonate,  potassium  carbonate.    Account  for  the  results. 

174.  Flame  tests.    Hold  a  platinum  wire  in  the  Bunsen 
flame  until  it  ceases  to  impart  any  color  to  the  flame ;  then 
dip  it  into  a  solution  of  a  compound  of  sodium  and  hold  it  in 
the  outer  film  of  the  base  of  a  Bunsen  flame.    Note  the  color 
imparted  to  the  flame.    Clean  the  wire  by  boiling  it  in  dilute 
hydrochloric  acid  and  repeat  the  flame  test  with  a  compound 
of  potassium. 

[116] 


175.  Ammonium  compounds,  a.  Recall  the  action  of  so- 
dium hydroxide  and  calcium  hydroxide  on  compounds  of 
ammonium.  Place  a  crystal  of  some  compound  of  ammo- 
nium in  a  small  beaker,  and  moisten  it  with  1  or  2  cc.  of  a 
solution  of  sodium  hydroxide.  Moisten  a  strip  of  red  litmus 
paper,  stick  it  on  the  bottom  of  a  watch  glass,  and  place  the 
latter  on  the  beaker  as  a  cover.  Why  does  the  litmus  paper 
turn  blue  ?  Will  this  serve  as  a  test  for  ammonium  salts  ? 

b.  Pour  2  or  3  cc.  of  a  solution  of  ferrous  sulfate  into  each 
of  two  test  tubes.   To  the  one  add  a  few  drops  of  an  aqueous 
solution  of  hydrogen  sulfide ;  to  the  other,  a  few  drops  of  a 
solution  of  ammonium  sulfide.    Account  for  the  results.   In 
preparing  sulfides  by  precipitation,  when  is  ammonium  sulfide 
used  in  place   of  hydrogen  sulfide  ?  How  is  it  prepared  ? 
What  changes  does  it  undergo  on  standing  exposed  to  air  ? 

c.  Ask  the  instructor  for  unknown  compounds  of  sodium, 
potassium,  or  ammonium,  and  see  if  you  can  identify  them. 
Recall  such  reactions  of  carbonates,  sulfates,  nitrates,   sul- 
fites,   sulfides,  chlorides,  bromides,  and  iodides  as  will  serve 
to  identify  them. 


[117] 


CHAPTER  XXIX 
THE  ALKALINE  EARTH  METALS 

176.  Calcium  hydroxide,    a.  Place    some    small   pieces  of 
marble  on  a  piece  of  wire  gauze  and  apply  a  strong  heat  for 
about  fifteen   minutes.    When    cool,   drop  the  residue  into 
25  cc.  of  water  and  stir.    Then  filter  the  liquid  and  divide  the 
filtrate  into  two  parts.    Blow  exhaled  air  through  one  portion 
(R).    Add  1  or  2  drops  of  ferric  chloride  to  the  other  por- 
tion (R).    What  other  hydroxides  would  be  precipitated  by 
calcium  hydroxide  ? 

b.  Place  about  10  cc.  of  clear  lime  water  in  a  test  tube  and 
conduct  carbon  dioxide  into  the  solution.  What  is  the  pre- 
cipitate (R)  ?  Continue  the  process  until  the  precipitate  dis- 
solves (R).  Divide  the  solution  into  two  parts.  Boil  one  part 
for  a  few  minutes  (R).  What  is  temporary  hardness  ?  Pre- 
pare a  dilute  solution  of  soap  and  shake  it  vigorously  in  a 
test  tube.  Does  it  form  a  froth  (lather)  ?  Add  to  it  some  of 
the  second  portion  of  the  solution  and  again  shake.  Does  it 
still  form  a  lather  ?  What  is  the  precipitate  ?  • 

177.  Calcium  chloride.    Dissolve  10  g.  of  marble  in  hydro- 
chloric acid   (R).    What    does    the   effervescence   indicate  ? 
Marble  is  likely  to  contain  a  little  carbonate  of  magnesium 
and  of  iron.    What  would  become  of  these  ?    Add  an  excess 
of  limewater  (R)  and  filter.    Evaporate  the  solution  to  a  vol- 
ume of  not  more  than  10  cc.  and  allow  it  to  crystallize.    What 
is  the  formula  of  the  crystals  ?    Drain  off  the  mother  liquid 
and  evaporate  it  to  dryness.    What  is  the  composition  of  the 
residue  ?    Expose  a  small  piece  of  it  to  the  air  for  an  hour 
and  account  for  the  results.    Dissolve  the  crystals  in  a  little 
water  and  divide  the  solution   into  two  portions.    To  one 

[118] 


portion  add  a  few  drops  of  ammonium  carbonate  (R) ;  to  the 
other  add  a  few  drops  of  solutions  of  ammonium  hydroxide 
and  disodium  phosphate  (R). 

178.  Calcium  sulfate.    a.  Heat  a  crystal  of  gypsum  in  a 
test  tube  (R).    Place  on  a  glass  plate  some  metal  object,  such 
as  a  file,  which  has  been  smeared  with  a  drop  of  oil.    Pour 
over  the  object  a  thick  paste  prepared  by  adding  water  to 
plaster  of  Paris.    Set  it  aside  until  it  hardens,  then  remove 
the  object  and  note  the  result.    What  causes  the  paste  to 
harden  ?    For  what  is  plaster  of  Paris  used  ? 

b.  Shake  1  g.  of  calcium  sulfate  with  10  cc.  of  water  in.  a 
test  tube,  filter,  and  divide  the  solution  into  two  parts.  Test 
one  part  for  the  presence  of  sulfates.  Is  calcium  sulfate 
soluble  in  water  ?  Shake  the  other  part  with  1  or  2  cc.  of 
soap  solution.  Does  a  froth  form  ?  How  could  such  a  solution 
be  softened  ?  Distinguish  between  temporary  and  permanent 
hardness  of  water.  To  what  is  each  due  ? 

179.  Barium  chloride.    Weigh  out  accurately  in  a  porcelain 
crucible  from  1  to  2  g.  of  small  crystals  of  barium  chloride. 
Place  the  lid  on  the  crucible  and  heat  the  crucible  gently  for 
a  few  minutes,  holding  the  burner  in  the  hand  and  moving 
it  about  so  as  to  apply  the  heat  uniformly.    Finally,  apply  a 
strong  heat  for  five  minutes.    When  the  crucible  is  cool,  weigh. 
The  residue  is  BaCl2.    From  your  results  determine  the  num- 
ber of  molecules  of  water  of  hydration  in  the  crystals. 

180.  Analytical   reactions.     Place   in   separate   test   tubes 
solutions  of  a  compound  of  each  of  the  following  elements: 
calcium,  barium,  strontium.    Test  the  effect  of  each  on  the 
flame  by  means  of  a  platinum  wire,  as  in  §  174.    Add  a  few 
drops  of  a  solution  of  potassium  dichromate  to  each.    Test 
the  solubility  of  the  precipitates  in  acetic  acid.    Add  an  equal 
volume  of  a  saturated  solution  of  calcium  sulfate  to  solutions 
of  each,  heat  to  boiling,  and  set  aside  until  cool.    Note  all 
the  results.    How  could  you  distinguish  between  compounds 
of  the  three  elements  ? 


[119] 


CHAPTER  XXX 
THE  MAGNESIUM  FAMILY 

181.  Magnesium  carbonate.    Place  2  or  3g.  of  magnesium 
carbonate  (preferably  the  mineral  magnesite)  in  a  hard-glass 
tube  fitted  with  a  delivery  tube,  and  gently  heat  in  the  Bun- 
sen  flame,  passing  the  gas  evolved  through  clear  limewater. 
What  would  you  judge  as  to  the  ease  of  decomposition  of 
magnesium  carbonate  in  comparison  with  that  of  calcium  car- 
bonate ?    Increase  the  heat,  and  when  the  evolution  of  gas 
becomes  slow,  cool  the  solid  product  and  shake  it  with  water, 
testing  the  reaction  toward  litmus  paper. 

182.  Magnesium  chloride,    a.  What  is  the  formula  of  crys- 
tallized magnesium  chloride  ?    Place  a  little  of  the  solid  in  a 
test  tube  and  heat  it  gently.    Is  water  given   off  easily  ? 
With  blue  litmus  paper  keep  testing  the  drops  which  con- 
dense on  the  sides  of  the  tube.    How  do  you  account  for  the 
reaction  (R)  ?    After  most  of  the  moisture  is  driven  off,  add 
water  to  the  residue.    Is  it  soluble  ?    What  is  it  ?    What  is 
the  industrial  importance  of  this  fact  ? 

b.  Add  a  few  drops  of  ammonium  hydroxide  to  a  solution 
of  magnesium  chloride  (R).  Repeat  the  experiment,  first 
diluting  the  magnesium  chloride  with  an  equal  volume  of  a 
solution  of  ammonium  chloride.  Does  a  precipitate  form  ? 
Since  there  are  many  magnesium  ions  present,  what  other  ion 
must  have  largely  disappeared?  Can  you  account  for  this 
by  mass  action?  To  the  solution  add  disodium  phosphate. 
The  precipitate  has  the  formula  MgNH4PO4  (R). 

183.  Weight  of  magnesium  sulfate  obtained  from  a  known 
weight  of  magnesium  oxide.    Weigh  accurately  a  small  evap- 
orating dish;  then  introduce  from  0.3  to  0.5  g.  of  magnesium 

[120] 


oxide  and  accurately  weigh  the  dish  and  contents.  Add  a 
little  more  than  enough  dilute  sulfuric  acid  to  dissolve  the 
oxide  (1  part  acid  to  3  parts  water),  and  carefully  evap- 
orate to  dryness,  completing  the  operation  under  the  hood. 
Moisten  the  powder  with  1  or  2  drops  of  the  acid  and  again 
evaporate,  finally  heating  the  product  to  a  low  red  heat. 
Allow  the  crucible  to  cool,  and  weigh  accurately  the  result- 
ing magnesium  sulfate.  Compare  your  results  with  the 
theoretical  results. 

184.  Zinc.  a.  Place  a  small  piece  of  zinc  on  charcoal  and 
heat  it  in  the  oxidiz ing-flame  of  the  blowpipe  (R).    What  is 
the  color  of  the  product  while  hot  ?    Does  its  color  change 
on  cooling? 

b.  Try  the  solubility  of  zinc  oxide  in  sodium  hydroxide. 
Could  a  film  of  oxide  remain  on  a  piece  of  zinc  in  a  solution 
of  this  reagent  ?    If  zinc  is  perfectly  clean,  how  would  you 
expect  it  to   act   with  water?     (See   electromotive   series.) 
Should  zinc  be  soluble  in  a  solution  of  sodium  hydroxide  ? 
Try  it. 

c.  Repeat  §  182,  &,  substituting  zinc  sulfate  for  magnesium 
chloride. 

185.  Combining  weight  of  zinc.    Repeat  §  183,  substitut- 
ing   for    the    magnesium    oxide    about    1  g.    of    pure    zinc, 
accurately    weighed,    and     for    the     sulfuric     acid,     dilute 
nitric   acid.    What  is  the  product  first   formed  ?    What  is 
obtained  on  heating  to  a  low  red  heat  ?    From  the  values 
obtained,   calculate  the   combining  weight  of   zinc  referred 
to   oxygen.     How   does    this    compare    with    the    combining 
weight  of  zinc  referred  to  hydrogen  as  unity  as  determined 
in  §  85  ? 

186.  Cadmium.    Obtain  about  5  cc.  of  a  solution  of  a  salt 
of  cadmium,  and  add  to  it  a  little  dilute  hydrochloric  acid. 
Then  pass  hydrogen  sulfide  into  the  solution  or  add  a  solu- 
tion of  the  reagent  to  it  (R).    What  is  the  color  of  the  pre- 
cipitate ?    Is  it  soluble  in  dilute  hydrochloric  acid  ?    Test  its 
solubility  in  concentrated  hydrochloric  acid  (?).    For  what  is 
it  used  ? 

[121] 


187.  Analytical  reactions.  Pour  into  separate  test  tubes  a 
solution  of  a  compound  of  magnesium,  zinc,  and  cadmium. 
Add  a  few  drops  of  hydrochloric  acid  to  each  solution  and 
then  pass  in  hydrogen  sulfide.  Note  the  result  (R).  What 
would  you  infer  as  to  the  solubility  of  the  sulfides  of  these 
metals  in  dilute  acids  ?  Add  ammonium  sulfide  to  separate 
solutions  of  compounds  of  magnesium  and  zinc.  Repeat, 
adding  an  equal  volume  of  ammonium  chloride  solution 
before  adding  the  ammonium  sulfide.  Explain  (R).  How 
could  you  detect  the  three  elements  in  the  presence  of 
each  other? 


[122] 


CHAPTER  XXXI 
ALUMINIUM 

188.  Aluminium  and  its  hydroxide,    a.  Dissolve  a  small 
piece  of  aluminium  in  hydrochloric  acid.     Where  must  the 
metal   stand  in  reference  to  hydrogen  in  the  electromotive 
series  ?    Add  sodium  hydroxide,  a  drop  at  a   time,  until  a 
precipitate  is  produced  (R).    Continue  the  addition  with  fre- 
quent stirring  (R).    When  solution  has  been  effected,  add 
hydrochloric   acid,   a  drop   at  a  time  (R).    By  what  name 
would  you  designate  a  hydroxide  with  such  properties  ? 

b.  Try  the  action  of  aluminium  on  boiling  water.    From 
the  place  of  the  metal  in  the  electromotive  series  would  you 
expect  it  to  decompose  water  ?     Now  try  the  action  of  the 
metal  on  a  solution  of  sodium  hydroxide  (R).    Does  the  fact 
that  the  hydroxide  is  soluble  in  alkalies  suggest  a  reason, 
for  the  fact  that,  while  the  metal  is  not  acted  upon  by  water, 
it  dissolves  in  alkalies  ?    Polish  the  surface  of  a  piece  of  alu- 
minium and  dip  it  into  a  solution  of  mercuric  chloride.  What 
becomes  of  the  mercury  (electromotive  series)  ?    The  mer- 
cury keeps  the  hydroxide  from  sticking  to  the   aluminium. 
Does  water  now  attack  the  metal  ? 

c.  Try  the  action  of  a  solution  of  ammonium  hydroxide 
on  a  solution  of  aluminium  chloride  (R).    Is  the  hydroxide 
dissolved  by  an  excess  of  ammonium  hydroxide  ? 

189.  Aluminium  salts,    a.  To  a  solution  of  aluminium  sul- 
fate  (or  any  other  soluble  salt  of  the  metal)  add  a  solution 
of  sodium  carbonate.    What  gas  is  evolved  ?    What  solid  is 
precipitated  ?     How  could  you  prove  it  is  not  a  carbonate  ? 
Repeat  the  experiment,  using  ammonium  sulfide  in  place  of 
sodium  carbonate  (R).   Pass  hydrogen  sulfide  into  a  solution 

[123] 


of  a  salt  of  aluminium.  Does  a  precipitate  form  ?  How  could 
you  test  for  zinc  and  aluminium  in  the  presence  of  each 
other  ?  How  test  for  magnesium  and  aluminium  ? 

b.  Obtain   about  1  g.   of   some  dry  compound   containing 
aluminium,  and  heat  it  on  charcoal  with  the  blowpipe  (R). 
Moisten   the  residue  with  a  drop  or  two  of  a  solution  of 
cobalt  nitrate  and  heat  it  once  more.    Note  the  color  of  the 
residue.    This  constitutes  a  test  for  compounds  containing 
aluminium. 

c.  Repeat  the  experiment,  substituting  a  compound  of  zinc 
for  that  of  aluminium.   The  green  product  is  called  Rinmann's 
green. 

d.  Mix  a  little  dry  alum  and  sodium  bicarbonate  and  rub 
them  together.    Pour  a  little  water  on  the  mixture  and  note 
the  result. 

190.  Aluminium   nitride.     Mix    thoroughly    10  g.    of    fine 
aluminium  powder  with  1  g.  of  lampblack.     Place  the  mix- 
ture in  the  form  of  a  cone  on  a  brick  or  iron  plate.    In  the 
top  of  the  cone  introduce  a  piece  of  magnesium  ribbon  about 
5  cm.  long.    Now  light  the  magnesium  ribbon.    The  combus- 
tion progresses  safely  and  without  explosion.    When  the  mass 
is  cooled,  note  the  crystals  of  aluminium  nitride  mixed  with 
crystals  of  aluminium  oxide.    Place  some  of  the  product  in 
a  test  tube  and  cover  it  with  a  solution  of  caustic  soda,  heat 
gently,  and  note  odor  of  gas  evolved  (?).  What  possible  use 
does  this  suggest  for  aluminium  ? 

191.  Double   salts,    a.  Alums.    What   is    a   double  salt? 
What  is  an  alum  ?    Calculate  the  weight  of  aluminium  sul- 
fate  (remember  that  it  is  a  hydrated  salt)  and  of  ammonium 
sulfate  required  for  the  preparation  of  25  g.  of  crystallized 
alum.    Dissolve  these  separately  in  hot  water  so  that  the  com- 
bined volume  will  be  about  75  cc.    Unite  the  hot  solutions 
and  set  the  product  aside  to  crystallize.    Can  you  make  out 
the  form  of  the  crystals  ?    Test  the  reaction  to  litmus  paper 
of  a  solution  of  a  few  pure  crystals. 

b.    Carnallite.    What  is  the  formula  of  carnallite  ?     Calcu- 
late the  weight  of  the  individual  salts  necessary  to  make  25  g. 

[124] 


of  crystallized  carnallite.  Which  of  these  are  hydrated  ? 
Weigh  out  the  required  amount  of  magnesium  chloride  and 
about  one  third  more  than  the  required  amount  of  potassium 
chloride,  mix  the  salts,  and  dissolve  them  in  the  least  possible 
volume  of  hot  water.  Allow  the  solution  to  cool,  decant  the 
mother  liquor,  and  wash  the  crystals  with  a  very  little  cold 
water.  See  if  you  can  detect  both  magnesium  and  potassium 
in  the  crystals.  Is  this  a  double  salt  or  a  complex  salt? 


[125] 


CHAPTER  XXXII 
THE  IRON  FAMILY 

192.  Reactions  of  the  ferrous  ion.     a.  Place  about  1  g.  of 
iron  in  a  beaker  and  cover  the  material  with  water.    Iron  by 
alcohol  dissolves  the  fastest,  but  tacks  or  clean  turnings  will 
do.    Add  dilute  hydrochloric  acid,  a  small  portion  at  a  time, 
so  as  to  keep  up  a  brisk  evolution  of  gas  (R).     Note  the 
odor.    It  is  chiefly  due  to  phosphine ;   how  do  you  account 
for  the  presence  of  this  substance  ?    Note  the  choking  effect 
of  the  gas  when  breathed.    Hold  a  nonluminous  flame  over 
the   beaker  for  a  moment.     How   do    you   account  for  the 
flashes  of  light  ?    Before  the  iron  has  all  dissolved,  filter  and 
at  once  make  the  following  tests,  using  1  or  2  cc.  for  each. 

b.  Try  the  action  of  a  solution  of  ammonium  sulfocyanate 
(NH4CNS).  Pass  hydrogen  sulfide  into  a  little  of  the  solu- 
tion (R).  Test  the  action  of  ammonium  sulfide  (R).  How 
do  you  account  for  the  difference  in  the  results  of  the  last 
two  experiments  ?  Add  sodium  hydroxide  to  some  of  the 
solution  (R).  What  is  the  color  of  the  precipitate  ?  What 
change  occurs  after  the  contents  of  the  test  tube  have  stood 
exposed  to  air  (R)  ? 

193.  Ferrous  ammonium  sulfate.    Dissolve  10  g.  of  iron  in 
dilute    sulfuric    acid   and   filter  from    the   insoluble    residue 
(what  is  it  ?).    What  weight  of  ferrous  sulfate  should  be  in 
the  filtrate  ?    Dilute  it  to  about  50  cc.    What  would  be  an 
equimolecular  weight  of  ammonium  sulfate  ?    Weigh  out  this 
amount  and  add  it  to  the  hot  solution  of  ferrous  sulfate.    Set 
the  beaker  aside  for  the  crystallization  of  the  salt.    Filter  off 
the  crystals,  wash  them  with  a  very  little  cold  water,  and 
spread  them  on  filter  paper  to  dry.    What  is  the  formula? 

[126] 


194.  Reactions  of  the  ferric  ion.    a.  As  in  §  192,  dissolve 
about  1  g.  of  iron  in  dilute  hydrochloric  acid  and  then  treat 
the  solution  with  3  or  4  cc.  of  aqua  regia.    (What  does  this 
produce  (R)  ?)     Use  portions  of  the  resulting  solution  for 
the  experiments  below. 

b.  To  one  portion  of  the  solution  obtained  in  a  add  am- 
monium hydroxide  in  excess  (R).     To  another  portion  add 
sodium  hydroxide  (R).    Is  the  product  soluble  in  an  excess 
of  the   reagent?     Test  a  third  portion  with  a  solution   of 
ammonium  sulfocyanate  (R).    To  a  fourth  add  sodium  car- 
bonate in  excess.    What  is  the  precipitate  (R)  ?    To  a  fifth 
portion  add  ammonium  sulfide   in  excess.     Is   the  product 
ferric  sulfide  (R)  ?    How  can  you  tell  ?    Dissolve  some  of 
the  precipitate  in  hydrochloric  acid.     The  milky  residue  is 
sulfur.    Can  you  account  for  it? 

c.  Boil  the  remainder  of  the  solution  prepared  in  a,  with 
the  addition  of  hydrochloric  acid,  until  the  aqua  regia  has 
been  destroyed.    Then  add  about  half  a  gram  of  zinc,  warm- 
ing the  whole  to  maintain  a  brisk   evolution   of  hydrogen. 
What  change  in  color  do  you  note  ?   From  time  to  time  test 
small  portions   with  ammonium   sulfocyanate,   adding    more 
zinc  if  necessary.    What  inference  do  you  draw? 

d.  If  you  dissolve  a  little  iron  in  nitric  acid,  would  you 
expect  to   obtain   a  ferric  salt  or  a  ferrous   salt?     Try  it. 
Obtain  a  crystal  of  ferric  nitrate,  put  it  in  a  test  tube,  and 
add  a  little  water.    How  do  you  explain  the  result?    Since 
nitric  acid  is  a  strong  acid,  how  would  you  describe  ferric 
hydroxide  ? 

195.  Ferric  ammonium  sulfate.    Dissolve  10  g.  of  crystals 
of  ferrous  sulfate  in  20  cc.  of  water.   Add  to  the  solution  the 
amount  of  sulfuric  acid  necessary  to  convert  the  salt  to  the 
ferric   state;  then  add  nitric  acid,   a  drop  at  a  time,  until 
the  color  no  longer  changes.    Evaporate  (hood)  to  a  sirupy 
mass  and  dissolve  in  a  little  hot  water.    Add  to  the  result- 
ing solution  the  amount  of  ammonium  sulfate  necessary  to 
form  ferric  ammonium  sulfate,  and  heat  gently  until  the  salt 
is  dissolved.    Set  the  resulting  solution  aside  to  crystallize. 

[127] 


Examine  the  form  of  the  crystals.  To  what  class  of  com- 
pounds does  ferric  ammonium  sulfate  belong  ?  Write  the 
equations  for  the  reactions  involved. 

196.  Potassium   ferrocyanide   and   potassium   ferricyanide. 
What  are  the  formulas  of  these  two  salts  ?    Does  either  of 
them  give  reactions  for  ferrous  or  for  ferric   ions  ?     What 
reactions  will  show  this  ?    Try  the  action  of  a  solution  of 
potassium  ferrocyanide  on  a  ferrous  salt  and  then  on  a  ferric 
salt   (R).    Make  a  similar  experiment,  using  a  solution  of 
potassium  ferricyanide  and  trying  its  action  on  a  ferric  and 
on  a  ferrous  salt  (R).    Tabulate  your  results. 

197.  Cobalt  and  nickel.    Test  separate  solutions  of  a  salt  of 
cobalt  and  a  salt  of  nickel  with  the  borax  bead ;  with  a  solu- 
tion of  sodium  hydroxide ;  with  ammonium  sulfide.    Note  the 
results  (R). 

198.  Detection  of  iron,  aluminium,  calcium,  and  magnesium 
in  the  presence  of  each  other,   a.  To  solutions  of  a  salt  of  iron 
(ferric)  and  of  aluminium,  respectively,  add  sodium  hydrox- 
ide, a  drop  at  a  time,  until  an  excess  has  been  added  (R). 
Note    that   the    aluminium   compound   at   first   precipitated 
redissolves  on  the  further  addition  of  sodium  hydroxide.    To 
the  resulting  solution  add  hydrochloric  acid  until  the  solu- 
tion becomes  acid,  then  make  it  slightly  alkaline  with  ammo- 
nium hydroxide  and  heat  to  boiling.    Note  the  result.    Filter 
off  the  precipitate  obtained  by  adding  sodium  hydroxide  to 
the  iron  compound,  dissolve  a  small  portion  of  it  in  dilute 
hydrochloric  acid,  and  add  a  drop  of  potassium  sulfocyanate 
(R).    How  can  you  detect  iron  and  aluminium  compounds 
in  the  presence  of  each  other? 

b.  Pour  into  separate  test  tubes  about  2  cc.  of  solutions  of 
separate  compounds   of  each  of  the  following  metals:   iron, 
aluminium,  calcium,  magnesium.    Add  to  each  an  equal  vol- 
ume of  ammonium  chloride  solution  and  then  a  few  drops 
of  ammonium  hydroxide.    Note  the  results.    How  can  you 
separate  iron  and  aluminium  from  calcium  and  magnesium  ? 

c.  To  a  solution  of  a  salt  of  calcium  and  a  solution  of  a 
salt  of   magnesium,  respectively,   add   an    equal  volume   of 

[128] 


ammonium  chloride  solution,  then  a  few  drops  of  ammonium 
carbonate  solution  (R).  How  would  you  verify  the  presence 
of  magnesium  ? 

d.  Prepare  an  outline  for  the  detection  of  compounds  of 
iron,  aluminium,   calcium,  and  magnesium  when  mixed  to- 
gether in   the    same    solution.     Present  the    outline  to  the 
instructor  for   criticism   and   approval. 

e.  Dissolve  about  1  g.  of  limestone  in  dilute  hydrochloric 
acid,  adding  2  or  3  drops  of  nitric  acid  to  oxidize  any  iron 
to  the  ferric  state.    What  does  the  effervescence  indicate  ? 
What  is  the  solution  likely  to  contain  ?   Filter  off  any  residue 
(silica)  and  test  the  filtrate  for  the  presence  of  iron,  alu- 
minium, calcium,  and  magnesium,  according  to  your  outline. 
If  the  amount  of  iron  and  aluminium  present  is  small,  do 
not  attempt  to  separate  them  from  each  other. 


[129] 


CHAPTER  XXXIII 
COPPER;  MERCURY;  SILVER 

199.  Cuprous  chloride.  Powder  5g.  of  cupric  sulfate  and 
place  it  in  a  large  test  tube  or  a  small  beaker.  Add  20  cc.  of 
concentrated  hydrochloric  acid  and  about  10  g.  of  copper, 
preferably  in  a  form  that  has  a  large  surface.  What  ions  are 
now  present  ?  What  action  should  metallic  copper  have  on 
the  Cu+  +  ion?  Boil  the  mixture  vigorously  for  about  ten 
minutes  (hood),  and  then  pour  it  into  a  beaker  containing 
at  least  200  cc.  of  water.  The  white  precipitate  is  cuprous 
chloride  (CuCl)  (R).  Decant  the  solution  and  wash  the  pre- 
cipitate once  with  water.  Boil  a  little  of  the  cuprous  chloride 
suspended  in  water  (R).  Treat  a  small  portion  with  a  solu- 
tion of  sodium  hydroxide  (R).  Warm  a  third  portion  with 
aqua  regia  (R). 

200o  Cuprous  iodide.  To  2  or  3  cc.  of  a  solution  of  cupric 
sulfate  add  about  1  cc.  of  a  solution  of  potassium  iodide  (R). 
How  do  you  account  for  the  separation  of  iodine  ?  How  can 
you  remove  the  latter  from  the  insoluble  cuprous  iodide  ? 
Which  is  the  more  easily  prepared,  cuprous  chloride  or 
cuprous  iodide  ? 

201.  Reactions  of  cupric  salts,  a.  Recall  the  action  of  sul- 
furic  and  of  nitric  acid  on  copper  (R).  Dip  a  nail  into  a 
solution  of  cupric  sulfate  (R).  What  other  metals  would  act 
like  iron  ? 

b.  To  a  cold  solution  of  cupric  sulfate  add  one  half  its 
volume  of  a  solution  of  sodium  hydroxide  (R).  Heat  the 
solution  to  boiling  and  account  for  the  change  in  the  color  of 
the  precipitate  (R).  Is  cupric  hydroxide  soluble  in  sodium 
hydroxide  ? 

[130] 


c.  Try   the   action   of   a  solution  of  hydrogen  sulfide  on 
cupric  sulfate  (R)  ;  of  a  solution  of  ammonium  sulfide  (R). 
Add  a  drop  of  ammonium  hydroxide  to  a  dilute  solution  of 
cupric  sulfate  (R).    Continue  to  add  the  ammonium  hydrox- 
ide, drop  by  drop,  until  the  precipitate  dissolves.    How  does 
the  color  of  the  solution  compare  with  that  of  the  original 
cupric   sulfate  ?    Can  there  be  any  considerable  number  of 
cupric  ions  in  the  solution  ?    Why  ? 

d.  To  5  cc.  of  a  solution  of  cupric  sulfate  add  1  cc.  of  hy- 
drochloric acid ;  then  add  an  equal  volume  of  a  solution  of 
hydrogen  sulfide  (R).    Repeat  with  a  solution  of  zinc  sulfate. 
Try  the  action  of  ammonium  sulfide  upon  zinc  sulfate.    How 
can  you  detect  copper  and  zinc  in  the  presence  of  each  other  ? 

e.  From  the  results  obtained  in  §  37  calculate  the  number 
of  molecules  of  water  of  hydration  present  in  cupric  sulfate. 

202.  Tetraminocuprisulfate    (Cu(NH3)4S04 •  H20).      Powder 
about  10  g.  of  cupric  sulfate  and  dissolve  it  in  20  cc.  of  con- 
centrated ammonium  hydroxide  and  10  cc.  of  water.    While 
vigorously  stirring  the  solution,  add,  drop  by  drop,  about  25  cc. 
of  alcohol.    Collect  the  purple-blue  crystalline  precipitate  on 
a  filter.    Can  you  prove  that  the  compound  contains  copper  ? 
that  it  contains  ammonia  ?  that  it  is  a  sulfate  ? 

203.  Analysis  of  brass.    To  detect  the  presence  of  copper 
and  zinc  in  brass,  place  0.5  g.  of  brass  in  an  evaporating-dish 
(hood)  and  dissolve  it  in  as  little  nitric  acid  as  possible.    Note 
the  color  of  the  solution.    What  does  it  indicate  ?    Evaporate 
the  solution  just  to  dryness,  add  5  cc.  of  hydrochloric  acid 
and  5  cc.  of  water,  and  warm  gently.    Transfer  the  clear  liquid 
to  a  beaker,  dilute  to  100  cc.  with  hot  water,  and  pass  a  slow 
current  of  hydrogen  sulfide  through  the  solution  as  long  as 
a  precipitate  forms.    What  is  the  precipitate  ?   Filter,  and  test 
the  filtrate  with  hydrogen  sulfide  in  order  to  be  certain  that 
the  precipitation  is  complete.    If  no  more  precipitate  forms, 
evaporate  the  filtrate  to  half  its  volume,  add  a  few  drops  of 
nitric  acid,  and  again  heat  to  boiling.    When  the  solution  is 
cool,  add  ammonium  hydroxide  until  it  is  alkaline,  then  warm 
gently  and  set  aside  for  a  few  minutes.    A  slight  precipitate 

[131] 


of  ferric  hydroxide  (Fe(OH)3)  may  form,  due  to  traces  of 
iron  in  the  brass.  Filter,  and  add  ammonium  sulfide  to  the 
clear  nitrate.  What  is  the  composition  of  the  precipitate? 

204.  Mercuric  compounds,  a.  Note  the  physical  properties 
of  mercury.    Place  in  a  small  beaker  a  globule  of  mercury  as 
large  as  a  grain  of  wheat,  and  add  (hood)  just  enough  con- 
centrated nitric  acid  to  dissolve  it.    Write  the  equation  for 
the  reaction  on  the  supposition  that  mercuric  nitrate  is  formed. 
Dilute  the  product  with  10  cc.  of  water  and  place  a  copper 
cent  in  the  solution.    After  a  few  minutes  remove  the  coin 
and  polish  it  with  a  piece  of  cloth.    Account  for  the  result. 

b.  For  what  purpose  have  we  used  mercuric  oxide  ?    Place 
0.5  g.  of  it  in  a  test  tube  and  dissolve  it  in  as  little  nitric 
acid  as  possible  (R),  then  add  water  until  the  test  tube  is 
one  fourth  full.    Divide  the  solution  into  three  equal  parts. 
To  one  part  add  a  little  hydrochloric  acid  (R);  to  a  second 
part  add  ammonium  hydroxide  (is  the  product  mercuric  hy- 
droxide ?    how  can  you  prove  it  ?)  ;    to  the  third  part  add  a 
small  piece  of  zinc  or  tin,  and  after  a  part  of  the  metal  has 
dissolved,  add  hydrochloric  acid  (R). 

c.  Prepare  a  little  mercuric  iodide  by  precipitation.    What 
is  its  color  ?  Collect  it  on  a  filter  paper  and  allow  it  to  dry. 
Scrape  a  little  of  the  powder  into  a  test  tube  and  very  carefully 
warm  it  at  some  distance  above  the  flame.    What  change  in 
color  takes  place  ?  Is  this  change  of  color  reversed  on  cooling  ? 

205.  Mercurous  compounds.    Put  a  small  globule  of  mer- 
cury in  a  test  tube  and  add  a  little  dilute  nitric  acid,  warm- 
ing   gently    and    taking    care    that    some    mercury   remains 
undissolved.    Divide  the  solution  into  three  parts.    To  one 
part  add  hydrochloric  acid  or  a  soluble  chloride  (R),  to  a 
second  add  ammonium  hydroxide   (R),   to  the  third  add  a 
little  concentrated  nitric  acid  (R).    When  action  has  ceased, 
add  1  or  2  cc.  of  hydrochloric  acid  (R).    How  can  you  dis- 
tinguish between  mercuric  and  mercurous  salts  ? 

206.  Compounds  of  silver,    a.    Test  the  reaction  of  a  solu- 
tion of  silver  nitrate  upon  litmus  paper.    What  would  you 
conclude  as  to  the  strength  of  silver  hydroxide  as  a  base  ? 

[132] 


Place  a  drop  of  silver  nitrate  solution  on  a  piece  of  cotton 
cloth  and  warm  gently.  Can  you  wash  the  stain  away?  What 
is  it?  Try  ammonia  water.  Owing  to  the  permanence  of  this 
stain,  silver  nitrate  is  used  in  making  indelible  ink. 

b.  To   5  cc.   of    a  solution  of   silver   nitrate   add    sodium 
hydroxide  in   excess    (R).    Wash  the  black   precipitate  re- 
peatedly with  hot  water.    Test  its  solubility  in  nitric  acid 
(R)  ;  in  ammonia  water  (R). 

c.  To  2  or  3  cc.  of  a  solution  of  silver  nitrate  add  hydro- 
chloric acid  (R).    Is  the  chloride  soluble  in  nitric  acid  ?  in 
ammonium  hydroxide  (R)  ?  in  sodium  thiosulfate  ?  Prepare 
sufficient  silver  sulfide  to  make  note  of  its  color  and  solu- 
bility in  hydrochloric  acid. 

d.  Prepare  small  amounts  of  the  chloride,  the  bromide,  and 
the  iodide  of  silver  (R).    Expose  to  sunlight  the  test  tubes 
containing  the  precipitates,  and  note  any  changes.    For  what 
are  these  compounds  used?    Prepare  sufficient  silver  sulfide 
to  note  the  color. 

207.  Preparation  of  pure  silver.  Place  2  or  3  g.  of  an  alloy 
of  copper  and  silver  in  a  small  beaker  and  add  (hood)  suffi- 
cient nitric  acid  to  dissolve  it.  The  solution  may  be  hastened 
by  applying  a  gentle  heat.  When  the  solution  is  complete, 
dilute  the  product  with  about  25  cc.  of  water.  Account  for 
the  color  of  the  liquid.  Now  add  a  solution  of  sodium  chlo- 
ride until  a  precipitate  ceases  to  form.  With  stirring,  the 
precipitate  (what  is  it  ? )  settles  to  the  bottom  of  the  beaker. 
Carefully  decant  the  clear,  supernatant  liquid  and  test  it  for 
the  presence  of  copper  (§  201,  c).  Wash  the  precipitate  two 
or  three  times  by  pouring  hot  water  over  it  and  decanting. 
Finally,  remove  any  remaining  water  by  filtration.  Mix  the 
moist  precipitate  with  an  equal  amount  of  sodium  carbonate, 
transfer  the  mixture  to  a  small  cavity  in  a  piece  of  charcoal, 
and  heat  it  with  a  blowpipe.  The  silver  salt  is  gradually 
reduced  to  metallic  silver,  which  will  fuse  into  a  globule  if 
sufficient  heat  is  applied. 


[133] 


CHAPTER  XXXIV 

TIN  AND  LEAD 

208.  Stannous  compounds,    a.  Dissolve  about  0.5  g.  of  tin 
in  hydrochloric  acid  (R).    Why  should  this  metal  dissolve 
so  much  more  slowly  than  zinc  ?   Cool,  dilute  to  10  cc.  with 
water,  and  use  for  the  experiments  below. 

b.  Place  about  2  cc.  of  a  solution  of  mercuric  chloride  in 
a  test  tube  and  add  stannous  chloride  drop  by  drop  (R). 
Does  the  color  of  the  precipitate  change  with  the  continued 
addition  of  the  reagent  (R)  ? 

c.  To  1  or  2  cc.  of  the  solution  add  a  little  aqua  regia, 
boil  until  the  acid  is  expelled,  and  repeat  the  experiment 
with  mercuric  chloride.     Why  does  no  precipitate  form? 

d.  Pour  a  few  drops  into  a  solution  of  hydrogen  sulfide  (R). 

e.  Add  a  few  drops  of  a  solution  of  gold  chloride  to  about 
5  cc.  of  water.    To  50  cc.  of  water  add  1  or  2  drops  of  the 
solution  of  stannous  chloride  and  1  or  2  drops  of  a  solution 
of  sodium  hydroxide.    Add  a  little  of  this  solution,  drop  by 
drop,  to  the  solution  of  gold  chloride.    A  deep  rose-purple 
color  slowly  develops,  known  as  the  purple  of  Cassius.    It  is 
due  to  finely  divided  (colloidal)  gold  (R). 

209.  Stannic  compounds,    a.  Add  a  few  drops  of  concen- 
trated nitric  acid  to  a  small  piece  of  tin,  heating  gently  if 
necessary.    What  is  the  compound  formed  ?    Is  it  soluble  in 
water  ?  in  hydrochloric  acid  ? 

b.  Dissolve  a  small  piece  of  tin  in  aqua  regia.  What  com- 
pound is  formed  (R)  ?  Pour  a  few  drops  of  the  solution  into 
a  solution  of  hydrogen  sulfide  (R).  Treat  a  little  of  the  so- 
lution with  a  few  drops  of  sodium  hydroxide  (R).  Is  the 
precipitate  soluble  in  excess  of  the  reagent? 

[134] 


210.  Crystallized  stannic  sulfide.     Mix  together  intimately 
5g.   of    stannous    sulfide,    2g.    of    ammonium    chloride,   and 
2.5  g.  of  sulfur,  placing  the  mixture  in  a  test  tube.     Clamp 
the  tube  loosely  in  a  nearly  horizontal  position   (hood)  and 
carefully  heat  it  with  a  wing-top  burner  held  in  the  hand, 
rotating  the  tube  by  twirling  the  rim  with  the  fingers,  and 
moving   the   burner   constantly.     Regulate  the  heat  so  that 
fumes  of  ammonium  chloride   are  steadily  evolved.     When 
nearly  all  of  the  fumes  have  escaped,  allow  the  tube  to  cool, 
and  tap  the  crystals  out  on  a  piece  of  paper.    To  what  ele- 
ment is  sulfur  similar  ?    Is  the  conversion  of  stannous  sulfide 
into  stannic  sulfide  a  process  of  oxidation  ? 

211.  Reactions  of  lead.    a.  Note  the  physical  properties  of 
the  metal.    Heat  a  small  piece  on  charcoal  in  the  oxidizing- 
flame.    Note  the  incrustation  formed  (R). 

b.  Place  1  g.  of  the  metal  in  an  evaporating-dish  (hood)  and 
add  20  cc.  of  water  and  5  cc.  of  nitric  acid.  Support  the  dish 
on  a  wire  gauze  and  heat  gently  until  the  metal  is  dissolved 
(R).  Evaporate  until  the  volume  is  reduced  about  one  half. 
What  is  the  composition  of  the  white  body  which  separates  ? 
Dilute  to  100  cc.  and  filter,  if  necessary,  to  obtain  a  clear  solu- 
tion. Now  test  small  portions  of  this  with  hydrogen  sulfide, 
sulfuric  acid,  potassium  chromate,  and  ammonium  carbonate 
respectively  (R).  Note  the  color  and  composition  of  the  pre- 
cipitates. Add  a  few  drops  of  hydrochloric  acid  to  a  test  tube 
one  fourth  full  of  the  solution  (R).  Heat  this  to  boiling,  and 
if  the  solution  does  not  become  clear,  add  just  enough  boiling 
water  to  dissolve  the  precipitate ;  then  set  it  aside  until  cool 
and  note  the  result.  How  can  you  distinguish  between  lead 
chloride  and  silver  chloride  ?  In  the  remainder  of  the  solution 
formed  by  dissolving  the  lead  in  nitric  acid,  suspend  a  piece  of 
zinc.  Set  aside  for  half  an  hour  and  note  the  result  (R).  Test 
the  solution  for  the  presence  of  zinc,  giving  the  method  used. 

212.  Oxides  of  lead.    Give  the  names  and  formulas  of  the 
oxides  of  lead.    Place  about  1  g.  of  lead  peroxide  in  each  of 
two  test  tubes.    To  the  one  add  a  few  drops  of  hydrochloric 
acid ;  to  the  other,  a  few  drops  of  sulfuric  acid.    Heat  each 

[135] 


gently  and  note  the  result  (R).  What  other  compound 
reacts  with  these  acids  in  a  similar  way  ?  Treat  a  little  red 
lead  with  dilute  nitric  acid  and  heat  gently.  After  the  action 
ceases,  dilute  and  filter.  Test  the  filtrate  for  lead.  Test  the 
residue  on  the  filter  paper  by  heating  with  hydrochloric  acid 
(R).  Try  the  action  of  hydrochloric  acid  and  nitric  acid  on 
litharge  (R). 

213.  The  storage  battery,  a.  Prepare  the  apparatus  illus- 
trated in  Fig.  55.  The  tube  A  is  the  large  test  tube  that 
was  used  in  the  preparation 
of  oxygen  ;  B  and  C  are  strips 
of  sheet  lead  rubbed  bright 
with  coarse  sandpaper.  These 
strips  should  be  about  1  cm. 
in  width,  and  should  be  bent 
at  right  angles  near  one  end 
so  as  to  rest  on  the  rim  of  the 
test  tube.  In  this  bent  end  a 
small  hole  is  made  by  driving 
a  wire  nail  through  the  lead, 
and  a  copper  wire  is  passed 
through  the  hole  and  twisted 
to  a  tight  connection.  When 
the  lead  plates  are  suspended 
in  the  tube  A^  they  must  be 
far  enough  apart  to  be  sepa- 
rated by  a  glass  tube  D,  so  as  to  prevent  short-circuiting. 
The  whole  cell  may  be  slipped  into  a  wide-mouthed  bottle 
as  a  support.  E  and  F  are  binding-posts,  terminals  of  two 
storage  cells  connected  in  series. 

b.  As  a  preliminary  experiment  obtain  1  or  2  cc.  of  a 
solution  of  starch  (side  shelf)  and  dissolve  in  it  a  crystal  of 
potassium  iodide.  Place  a  clean  sheet  of  filter  paper  on  a 
clean  and  dry  spot  on  the  table  close  to  the  binding-posts 
E  and  F,  and  on  this  pour  a  few  drops  of  the  solution  just 
prepared.  Attach  two  wires  to  the  binding-posts  and  bring 
the  two  free  ends  upon  the  moist  starch  paper  about  1  cm. 

[136] 


j, 


apart.  Is  there  any  change  of  color  where  the  wires  touch 
the  paper  ?  At  which  pole  does  it  occur  ?  How  do  you  ex- 
plain it  (R)  ?  Will  this  reaction  serve  to  detect  an  electric 
current  ? 

c.  Prepare  an  electrolyte  for  the  cell  by  mixing  equal 
volumes  of  concentrated  sulfuric  acid  and  water,  making 
enough  of  the  solution  to  fill  the  tube  A  nearly  full.  When 
the  solution  has  cooled,  pour  it  into  the  tube  A  and  bring 
the  ends  of  the  two  wires  connected  with  B  and  C  upon  the 
starch  paper.  Can  you  detect  any  current  ?  Is  your  cell  act- 
ing as  a  battery  ?  Now  connect  the  wires  with  the  binding- 
posts.  What  changes  in  the  cell  do  you  note  ?  Which  plate 
turns  brown  ?  What  is  the  brown  substance  ?  What  is 
formed  at  the  other  plate  ?  What  gases  are  evolved  ?  When 

Jr  o 

one  plate  has  become  decidedly  brown,  disconnect  the  wires 
from  E  and  F  and  bring  the  free  ends  upon  the  starch  paper, 
assuring  yourself  that  it  is  still  moist.  Is  iodine  set  free  now  ? 
What  is  the  source  of  the  current  (R)  ? 

214.  Analysis  of  solder.  What  two  metals  are  present  in 
ordinary  solder  ?  Devise  a  method  for  detecting  their  pres- 
ence in  it  and  submit  it  to  the  instructor  for  approval.  When 
your  method  has  been  approved,  try  it. 


[137J 


CHAPTER   XXXV 
MANGANESE  AND  CHROMIUM 

215.  Salts  of  manganese.     To  a  solution    of   manganous 
chloride  add  a  solution  of  hydrogen  sulfide  (R);   a  solution 
of  ammonium  sulfide  (R).    Try  the  action  of  sodium  hydrox- 
ide and  of   ammonium  hydroxide   (R).     Is  the   precipitate 
soluble  in  excess  of  the  reagents  ?  Does  it  undergo  any  change 
on  exposure  to  the  air  (R)  ? 

216.  Manganates  and  permanganates,    a.  In  a  mortar  grind 
5  g.  each  of  manganese  dioxide  and  potassium  hydroxide  and 
2.5  g.  of  potassium  chlorate.    Transfer  the  mixture  to  a  small 
iron  dish  and  heat  to  redness  until  the  mass  fuses  and  becomes 
deep  green.   The  green  material  has  the  composition  K2MnO4 
(R).   When  the  mass  has  cooled,  add  cold  water  and  dissolve 
as  much  of  it  as  possible,  decanting  the  solution  from  the 
residue. 

b.  The  manganate  is  stable  only  in  the  presence  of  an 
excess  of  alkali.    What  would  you  argue  as  to  the  strength 
of  manganic  acid  ?    In  water  it  decomposes  according  to  the 
equation 

3  K2MnO4  +  4  H2O  =  2  KMnO4  +  Mn(OH)4  +  4  KOH 

Add  dilute  nitric  acid  to  the  solution.  What  change  in  color 
do  you  note  ?  The  compound  KMnO4  is  deep  purple  in  color. 
Does  a  precipitate  form  ?  What  is  it  ?  Remembering  that 
the  reaction  is  reversible,  what  is  the  reason  that  nitric  acid 
promotes  the  reaction  from  left  to  right  ?  How  would  you 
promote  the  reaction  from  right  to  left  ?  Try  it. 

c.  In  the  presence  of  sulfuric  acid,  how  does  potassium 
permanganate  act  as  an  oxidizing  agent  (R)  ?   What  becomes 

[138] 


of  the  manganese  ?  Are  these  compounds  notably  colored  ? 
Add  a  little  acidified  solution  of  permanganate  to  a  solution 
of  ferrous  sulfate.  How  do  you  account  for  the  disappearance 
of  the  color  (R)?  In  a  similar  way  try  the  action  of  per- 
manganate upon  oxalic  acid  and  sulfurous  acid  (R).  Pour  a 
little  hydrochloric  acid  upon  two  or  three  crystals  of  potassium 
permanganate  (R). 

d.  Make  a  solution  of  potassium  permanganate  by  adding 
1  g.  of  the  salt  to  100  cc.  of  water,  and  fill  a  burette  with  the 
solution.  Weigh  accurately  0.5  g.  of  pure  iron  wire  and  dis- 
solve it  in  an  excess  of  dilute  sulfuric  acid.  What  is  present 
in  the  solution  ?  Slowly  add  the  permanganate  from  the 
burette,  constantly  stirring  the  solution.  When  the  color  of 
a  drop  added  fades  rather  slowly,  add  the  solution  more  care- 
fully, stopping  the  instant  a  permanent  pink  color  is  pro- 
duced (R).  The  continuance  of  the  color  shows  that  all  of 
the  ferrous  sulfate  has  been  oxidized.  From  the  weight  of 
the  iron  oxidized,  calculate  the  weight  of  permanganate  used 
up.  From  the  volume  of  the  permanganate,  calculate  the 
quantity  of  the  permanganate  in  a  liter  of  the  solution. 
Knowing  the  strength  of  the  solution,  could  you  determine 
the  amount  of  iron  present  in  a  sample  of  unknown  composi- 
tion ?  This  is  a  favorite  method  of  analysis. 

217.  Salts  of  chromium.    Try  the  effect  of  the  following 
reagents  upon  a  solution  of  a  chromium  salt,  writing  all  the 
equations:   ammonium  sulfide,  sodium  carbonate,  ammonium 
hydroxide.    Does    sodium   hydroxide    occasion  a  precipitate 
(R)  ?    Is  this  soluble  in  excess  of  the  reagent  (R)  ?    Try 
boiling  the   solution. 

218.  Chromates  and  dichromates.    a.  In  an  iron  dish  melt 
equal  weights  of  potassium  hydroxide  and  potassium  nitrate 
(about   5  g.  of   each),   and   gradually  stir  into  the  mixture 
about  5  g.  of  finely  powdered  chrome  iron   ore  or   3  g.  of 
chromic   oxide   (R).    The  nitrate  is  used    as    an    oxidizing 
agent.    Allow  the  fused  mass  to  cool,  and  extract  the  soluble 
portion  with  hot  water.     What  is  the  color?    To  what  is  it 
due  ?    To  3  or  4  cc.  of  the  solution  add  nitric  or  sulfuric  acid. 

[139] 


What  change  in  color  do  you  note  (R)  ?  Add  potassium 
hydroxide  to  this  solution.  Is  the  color  change  reversed  (R)  ? 
Write  the  equation  for  the  transformation  of  a  chromate  to  a 
dichromate  as  a  reversible  reaction.  Why  does  the  addition 
of  acid  throw  it  one  way,  while  bases  throw  it  the  other  ? 

b.  To  a  solution  of  a  chromate  add  a  solution  of  a  soluble 
salt  of  lead  (R).   Repeat,  using  a  salt  of  barium  (R).   Repeat 
these  experiments,  using  a  solution  of  a  dichromate  instead 
of  a  solution  of  a  chromate  (R). 

c.  Pour  2  or  3  cc.  of  concentrated  hydrochloric  acid  upon 
a  few  crystals  of  a  chromate  (R).    Repeat,  using  potassium 
dichromate  (R). 

d.  Place  three  or  four  crystals  of  potassium  dichromate  in 
a  test  tube  and  add  2  cc.  of  concentrated  sulfuric  acid.    The 
deep-red  crystals  are  chromic  anhydride  (CrO3)  (R).    Pour 
the   mixture    into    cold  water.     Do   the    crystals   dissolve  ? 
What   do   they  form   (R)  ? 


[140] 


APPENDIX 


DATA  FOR  CALCULATION  OF  GAS  VOLUMES 

The  relation  between  the  volume  of  a  gas  under  standard 
conditions,  Vg,  and  its  volume  V  when  measured  over  water  at 
a  pressure  P,  an  absolute  temperature  T,  and  an  aqueous  tension 
a  is  expressed  by  the  equation 

(P  -  a)  x  V  X  273 

760  x  T 

In  comparing  the  pressure  exerted  by  a  column  of  water  with 
that  exerted  by  a  column  of  mercury,  it  must  be  remembered  that 
the  density  of  mercury  is  13.56. 


TENSION  OF  AQUEOUS  VAPOR  AT  VARIOUS  TEMPERATURES 
EXPRESSED  IN  MILLIMETERS  OF  MERCURY 


TEMPERA- 


TURE 


10° 
11° 

12° 
13° 
14° 
15° 
16° 
17° 


PRES-       TEMPERA- 


SURE 

8.61 

9.20 

9.84 

10.51 

11.23 

11.98 

12.78 

13.62 

14.52 


PRES-       TEMPERA- 


TURE 


SURE 


19° 16.56 

20° 17.51 

21° 18.62 

22° 19.79 

23° 21.02 

24° 22.32 

25° 23.69 

26° 25.13 

27°  .  26.65 


TURE 


PRES- 
SURE 

29° 29.94 

30° 31.74 

31° 33.57 

32° 35.53 

33° 37.59 

34° 39.75 

35° 42.02 

36° 44.40 

37°  .  .  46.90 


18°  ,    15.46       28C 


28.25       100°   .  .  760.00 


WEIGHT   IN   GRAMS   OF    1    LITER  OF  VARIOUS  GASES   UNDER 
STANDARD  CONDITIONS,  AND  BOILING  POINTS  UNDER  PRES- 
SURE OF   760  MILLIMETERS 


WEIGHT 
NAME  OF  1  LITER 

Acetylene      .    .    .  1.1621 

Air 1.2928 

Ammonia  ....  0.7708 

Argon 1.7809 

Carbon  dioxide  .  1.9768 
Carbon  monoxide  .  1.2504 
Chlorine  ....  3.1674 

Helium 0.1782 

Hydrogen  ....  0.08987 


BOILING 
POINT 

-83.8° 


-33.5° 
-186.0° 

-78.2° 
-190.0° 

-33.6° 
-268.7° 
-252.7° 


WEIGHT  BOILING 

NAME             OF  1  LITER  POINT 

Hydrogen  chloride  .  1.6398  -83.1° 

Hydrogen  fluoride  .  0.893  + 19.4° 

Hydrogen  sulfide    .1.5392  —61.6° 

Methane      .    .    .    .0.7168  -164.0° 

Nitric  oxide    .    .    .1.3402  —153.0° 

Nitrogen      .    .    .    .1.2507  —195.7° 

Nitrous  oxide  .    .    .1.9777  -89.8° 

Oxygen 1.4290  -183.0° 

Sulfur  dioxide    ,     .  2.9266  -  8.0° 


[141] 


ELECTROCHEMICAL  SERIES  OF  METALS 


1.  Potassium. 

2.  Sodium. 

3.  Magnesium. 

4.  Aluminium. 


5.  Zinc. 

6.  Iron. 

7.  Tin. 

8.  Lead. 


9.  Hydrogen. 

10.  Copper. 

11.  Antimony. 

12.  Bismuth. 


13.  Mercury. 

14.  Silver. 

15.  Platinum. 

16.  Gold. 


OUTFIT  FOR  EACH  STUDENT  LOCKER 
A.   ARTICLES  RETURNABLE 


2  Bunsen  burners. 

1  burner,  wing-top. 

1  burette  clamp. 

1  screw  clamp. 

1  set  cork  borers,  Nos.  1-3. 

1  hemispherical  iron  dish,  75-mm. 

1  lead  dish. 

1  pair  forceps. 

1  pinchcock. 

1  deflagrating-spoon. 

1  test-tube  holder. 

1  test-tube  rack. 

1  mortar  and  pestle,  75-mm. 

1  piece  platinum  wire. 

1  porcelain  spatula,  130-mm. 

1  casserole,  75-mm. 

1  porcelain  crucible  and  lid,  No.  00. 

1  porcelain  dish,  No.  0, 

2  beakers,  60-cc. 
2  beakers,  100-cc. 
2  beakers,  150-cc. 


2  beakers,  225-cc. 
2  beakers,  325-cc. 
1  bottle,  narrow-mouthed,  1000-cc. 

1  bottle,  wide-mouthed,  1000-cc. 
4  bottles,  wide-mouthed,  250-cc. 

2  bottles,  wide-mouthed,  60-cc. 
1  drying-tube,  straight. 

1  flask,  Erlenmeyer,  120-cc. 
1  flask,  Florence,  120-cc. 
1  flask,  Florence,  250-cc. 

1  flask,  Florence,  500-cc. 

2  funnels,  65-mm. 

3  glass  plates,  75-mm.  x  75-mm. 
12  test  tubes,  130-mm. 

2  test  tubes,  hard-glass,  150-mm. 

1  test  tube,  graduated,  30-cc. 

1  funnel  tube. 

1  glass  retort,  200  cc. 

1  U-tube,  Marchand,  100-mm. 

1  watch  glass,  78-mm. 

1  Sargent  lock  and  2  keys.* 


B,   ARTICLES  NOT  RETURNABLE 


1  towel. 

1  clay  triangle. 

1  test-tube  brush. 

1  sponge. 

50  splints. 

1  box  matches. 

2  bottles  litmus  paper. 

1  hard-glass  tube,  500-mm.  x  7-mm. 
1  hard-glass  tube,  300-mm.  x  10-mm. 

1  pkg.  filter  paper,  110-mm.,  No.  595. 

2  oz.  of  glass  tubing,  6-mm, 


1  oz.  of  light-weight  glass  rod. 

1  cake  soap. 

3  ft.  of  rubber  tubing,  4-mm.,  No.  22. 
6  ft.  of  rubber  tubing,  6-mm.,  No.  22. 

2  rubber  stoppers,  two-hole,  No.  6. 
1  rubber  stopper,  one-hole,  No.  5. 
1  rubber  stopper,  one-hole,  No.  4. 
1  file,  triangular. 

1  file,  round. 
1  wire  gauze. 


[142] 


APPARATUS  TO  BE  AVAILABLE  WHEN  CALLED  FOR 

Burette,  50-cc 1  for  every  8  students 

Cylinder,  graduated,  1 50-cc 1  for  every  8  students 

Cylinder,  plain,  2500-cc.  (for  collecting  gases  in  tubes)  1  for  every  8  students 

Condenser,  Liebig,  300-mm 1  for  every  4  students 

Conductivity  apparatus  (see  note  to  §  72,  p.  148)     .  1  for  every  6  students 

Gas  tube,  graduated,  50-cc 1  for  every  4  students 

Gas  tube,  plain  (bomb  tube) 1  for  every  4  students 

Magnet,  horseshoe 1  for  every  10  students 

Reading  lens 1  for  every  10  students 

Separatory  funnel 1  for  every  4  students 

Thermometer,  ordinary  centigrade 1  for  every  4  students 

Thermometer  graduated  in  tenths  of  a  degree  (10°  — 

30°)  (see  note  to  §  131,  p.  150) 1  for  every  8  students 

Victor  Meyer  apparatus 2  or  3  sets 


APPARATUS  PROVIDED  FOR  GENERAL  USE 

Balance  and  weights  (see  note  to  §  7,  p.  148). 

Barometer,  mercurial. 

Blast  lamp. 

Pneumatic  trough  (see  note  to  §12,  p.  148).  If  locker  space  permits,  the 
pneumatic  trough  should  form  a  part  of  each  outfit. 

Ring  stands.  If  movable  ring  stands  are  used,  one  should  be  in  each  outfit. 
It  is  better  to  provide  the  laboratory  with  permanent  ring  stands  on 
each  desk.  These  are  easily  made  by  screwing  an  iron  rod  into  a 
flat  disk  that  in  turn  is  screwed  to  the  desk  top.  Local  workmen  can 
make  them. 


LIST  OF  CHEMICALS  FOR  A  CLASS  OF  TEN  STUDENTS 

In  the  following  list  will  be  found  an  estimate  of  the  actual 
quantities  of  reagents  required  for  a  class  of  ten  students.  In 
many  cases  these  quantities  are  less  than  the  smallest  commercial 
package,  and  no  supply  house  would  want  to  furnish  such  small 
quantities.  In  the  second  column  will  be  found  the  smallest  com- 
mercial package  that  will  be  sufficient,  the  figures  having  been 
supplied  through  the  kindness  of  the  Kauffman-Lattimer  Co. 
Columbus,  Ohio.  It  will  be  noticed  that  in  ordering  for  classes 
of  large  size  it  will  not  in  all  cases  be  necessary  to  multiply  the 
number  of  commercial  packages  by  the  number  of  pupils  in 
the  class. 

[143] 


ACTUAL 

QUANTITIES  COMMERCIAL 
IN  GRAMS      PACKAGES 

Acid,  acetic  (36  per  cent) 110  \  Ib. 

Acid,  formic  (90  per  cent) 500  1  Ib. 

Acid,  hydrochloric  (sp.  gr.  1.19) 2000  6  Ib. 

Acid,  nitric  (sp.  gr.  1.42) 300  7  Ib. 

Acid,  oxalic 80  |  Ib. 

Acid,  pyrogallic 40  1  oz. 

Acid,  sulf uric  (sp.  gr.  1.84) 2000  9  Ib. 

Agar f 10  1  oz 

Alcohol 500  1  qt. 

Aluminium  chloride 20  1  oz. 

Aluminium  foil 10  1  oz. 

Aluminium  sulfate 30  \  Ib. 

Ammonium  hydroxide  (sp.  gr.  0.90) 1000  4  Ib. 

Ammonium  carbonate 100  \  Ib. 

Ammonium  chloride 500  1  Ib. 

Ammonium  molybdate 5  1  oz. 

Ammonium  nitrate       80  ^  Ib. 

Ammonium  oxalate 20  1  oz. 

Ammonium  sulfate       100  \  Ib. 

Ammonium  sulfocyanate ' 10  1  oz. 

Antimony 50  2  oz. 

Antimony  oxide 10  1  oz. 

Antimony  sulfate 10  1  oz. 

Arsenic 10  1  oz. 

Arsenic  trioxide 10  \  Ib. 

Barium  chloride 30  ^  Ib. 

Barium  hydrate 10  1  oz. 

Barium  nitrate 30  ^  Ib. 

Barium  peroxide 50  2  oz. 

Benzene 500  1  pt. 

.Bismuth 40  2  oz. 

Bismuth  chloride 10  \  oz. 

Bismuth  nitrate 30  1  oz. 

Bone  black 30  \  Ib. 

Borax 100  \  Ib. 

Brass  turnings 10  1  oz. 

Bromine 40  2oz. 

Cadmium  sulfate 25  1  oz. 

Calcium  carbide 10  \  Ib. 

Calcium  chloride 500  1  Ib. 

Calcium  fluoride  (fluorite)   .    .    . 40  J  Ib. 

Calcium  nitrate 20  1  oz. 

Carbon  disulfide 80  \  Ib. 

Charcoal,  powdered 40  ^  Ib. 

[144] 


ACTUAL 

QUANTITIES  COMMERCIAL 
IN  GRAMS      PACKAGES 


Charcoal,  sticks 

Chloroform  .    .    .    .  ' 100 

Chromic  oxide 30 

Cobalt  nitrate 30 

Cochineal 

Copper  foil 50 

Copper  powder 170 

Copper  turnings 250 

Cotton       

Cottonseed  oil • 50 

Cupric  bromide 15 

Cupric  chloride 10 

Cupric  nitrate       30 

Cupric  oxide 140 

Cupric  sulfate       250 

Cuprous  oxide      30 

Ferric  chloride 20 

Ferric  oxide 10 

Ferrous  sulfate 120 

Ferrous  sulfide 250 

Glucose 1000 

Hydrogen  peroxide      50 

Iodine 15 

Iron  (by  alcohol) '.  170 

Lead  acetate 30 

Lead  nitrate 30 

Lead  oxide  (litharge) 30 

Lead  peroxide 30 

Lime 1000 

Limestone 60 

Litmus  cubes 3 

Magnesite 50 

Magnesium  chloride 180 

Magnesium  oxide 10 

Magnesium  powder      30 

Magnesium  ribbon 15 

Magnesium  sulfate 50 

Manganese  chloride 10 

Manganese  dioxide  (pyrolusite) 200 

Marble 200 

Mercuric  chloride 35 

Mercuric  oxide 10 

Mercuric  sulfate 10 

Mercury 1000 

[145] 


1  doz. 
*lb. 

loz. 
loz. 
1  oz. 
2oz. 
ilb. 

|lb. 

1  oz. 

*lb. 

ioz. 
loz. 
1  oz. 

ilb. 
ilb. 

loz. 

ilb. 

loz. 
ilb. 
1  Ib. 
21b. 


ilb. 
ilb. 
ilb. 
loz. 
2  Ib. 
lib. 
loz. 
ilb. 
ilb. 
2oz. 
1  oz. 
loz. 
1  Ib. 
loz. 
ilb. 
1  Ib. 
loz. 
loz. 
1  oz. 
21b. 


ACTUAL, 

QUANTITIES  COMMERCIAL 
J>T  GRAMS      PACKAGES 

Nickel  nitrate 20  1  oz. 

Paraffin 100  1  lb. 

Phenolphthalein i  oz. 

Phosphorus  (red) 50  2  oz. 

Phosphorus  (white) 50  2  oz. 

Phosphorus  pentoxide 10  1  oz. 

Phosphorus  trichloride 15  1  Oz. 

Potassium  alum 30  1  lb. 

Potassium  bromide 30  1  oz. 

Potassium  carbonate 20  2  oz. 

Potassium  chlorate 225  i  lb. 

Potassium  chloride 225  \  lb. 

Potassium  dichromate 100  \  lb. 

Potassium  ferricyanide 10  1  oz. 

Potassium  ferrocyanide 10  1  oz. 

Potassium  hydroxide 250  \  lb. 

Potassium  iodide 35  1  oz. 

Potassium  nitrate 280  \  lb. 

Potassium  perchlorate 20  1  oz. 

Potassium  permanganate 475  1  lb. 

Potassium  sulfate 10  \  lb. 

Rochelle  salts 200  \  lb. 

Silver  nitrate 30  1  oz. 

Soda  lime 200  \  lb. 

Sodium  (metal) 30  1  oz. 

Sodium  acetate 270  \  lb. 

Sodium  ammonium  phosphate 20  1  oz. 

Sodium  bicarbonate 30  1  lb. 

Sodium  bromide 20  1  oz. 

Sodium  carbonate 200  1  lb. 

Sodium  hydroxide 400  1  lb. 

Sodium  iodide      30  1  oz. 

Sodium  nitrate 250  \  lb. 

Sodium  nitrite 60  2  oz. 

Sodium  peroxide 50  2  oz. 

Sodium  phosphate,  disodium 30  \  lb. 

Sodium  phosphate,  trisodium 10  1  oz. 

Sodium  sulfate 300  1  lb. 

Sodium  sulfite 130  1  lb. 

Sodium  thiosulfate 30  1  lb. 

Stannous  sulfide 50  2  oz. 

Sulfur,  powdered 100  1  lb. 

Sulfur  roll 100  1  lb, 

Tartar  emetic 10  1  oz. 

[146] 


ACTUAL 

QUANTITIES  COMMERCIAL, 
IN  GKAMS      PACKAGES 

Tin,  mossy 250  £  Ib. 

Zinc,  granulated 170  |  Ib. 

Zinc,  mossy 500  1  Ib. 

Zinc  oxide 10  ^  Ib. 

Zinc  sulfate 40  1  Ib. 

MISCELLANEOUS  SUPPLIES 

Candles,  beeswax  (to  be  cut  into  pieces) 2 

Coal,  powdered 1  Ib. 

Coal  oil 1  qt. 

Corks,  Nos.  3,  6,  8,  10,  22,  24 2  doz.  each 

Emery  paper,  No.  0 2  sheets 

Glazed  paper 6  sheets 

Iron  nails  or  tacks 2  Ib. 

Iron  wire  (picture-frame) 1  box 

Labels,  Dennison's,  No.  219 2  boxes 

Lead  strips  (see  note  to  §  213,  p.  150) 10 

Olive  oil  (for  burns)      1  pt. 

Sand 2  Ib. 

Sandpaper,  No.  0       4  sheets 

Sawdust  (preferably  hardwood) 1  Ib. 

Starch .1  Ib. 

Sugar 1  Ib. 

Yeast 1  cake 

LABORATORY  SOLUTIONS 

Each,  working  desk  should  be  supplied  with  the  common  acids 
and  alkalies,  both  concentrated  and  dilute.  In  preparing  the 
dilute  reagents,  concentrated  sulfuric  acid  should  be  diluted  in 
the  ratio  1:4;  concentrated  hydrochloric  acid,  nitric  acid,  arid 
ammonia,  in  the  ratio  1  :  3. 

In  preparing  other  solutions  for  the  side  shelf  it  is  customary 
in  most  laboratories  to  make  10  per  cent  solutions,  though  it  is 
much  more  satisfactory  to  make  the  solutions  approximately 
normal  in  anhydrous  salt.  In  the  case  of  sparingly  soluble  salts, 
saturated  solutions  are  used.  With  expensive  reagents,  such  as 
silver  nitrate  or  potassium  iodide,  a  concentration  of  about  20  g. 
to  the  liter  is  sufficient  for  most  uses. 

It  is  convenient  to  have  the  bottles  on  the  side  shelf  arranged 
on  the  principle  of  a  wash  bottle,  a  rubber  tube  being  attached  to 

[147] 


the  glass  tube  into  which  one  blows.  Each  student  can  provide 
himself  with  a  glass  mouthpiece  to  insert  into  this  rubber  tube. 
If  the  delivery  tube  is  drawn  to  a  rather  fine  jet,  the  delivery 
will  be  slow  and  the  student  will  be  more  apt  to  heed  the  direc- 
tions to  use  small  quantities  of  reagents.  One  great  advantage 
of  this  arrangement  is  that  it  solves  the  problem  of  interchanged 
stoppers  and  spilled  reagents. 

COMMENTS  ON  SPECIAL  APPARATUS  AND  REAGENTS 

§  7.  The  balance.  It  is  very  desirable  that  the  balance  used  in 
quantitative  exercises  should  weigh  with  accuracy  to  1  mg.,  and 
such  a  balance  of  simple  type  may  be  bought  from  any  one  of  a 
number  of  firms.  Very  fair  results  may  be  secured  with  a  balance 
weighing  to  1  eg.,  and  such  a  balance  may  be  bought  for  about 
$9.  A  good  example  of  such  a  balance  is  the  Trcemner  No.  13. 
The  type  known  as  the  Harvard  balance  has  the  advantage  that 
the  weights  are  not  easily  lost. 

§  12.  Pneumatic  trough.  A  convenient  size  for  a  pneumatic 
trough  is  30  cm.  x  20  cm.  x  15  cm.  It  should  be  made  of  galva- 
nized sheet  iron  and  should  be  painted  to  protect  it  from  rust. 
A  supply  can  be  made  by  a  local  tinner.  Of  course  any  kind  of 
rather  deep  pan  will  serve  most  uses. 

§  15.  The  hard-glass  test  tube  should  be  about  15  cm.  in  length 
by  about  2  cm.  in  internal  diameter.  Such  tubes  can  be  used  in 
many  experiments  in  which  more  costly  retorts  and  flasks  are 
described,  and  a  good  supply  should  be  on  hand. 

§  39.  Starch  paste.  Starch  paste  for  the  side  shelf  may  be  pre- 
pared by  grinding  about  1  g.  of  starch  in  a  mortar  with  a  little 
cold  water.  The  thin  paste  is  then  poured  into  about  a  liter  of 
boiling  water,  boiled  for  one  minute,  and  allowed  to  cool. 

§  72.  Conductivity  apparatus.  The  apparatus  for  measuring 
the  conducting  power  of  solutions  (Fig.  36)  may  be  obtained  from 
dealers  in  apparatus  or  may  be  made  by  anyone  who  wires  houses 
for  electric  lights.  The  drawing  supplies  the  needed  information. 

§  74.  Equimolecular  acids.  It  is  a  tedious  task  to  make  solutions 
of  acids  of  exactly  standard  strength.  The  following  directions 
will  serve  nearly  all  the  purposes : 

[148] 


278  cc.  sulfuric  acid  (sp.  gr.  1.84)  will  make  5  liters  of  2  N  acid. 
629  cc.  nitric  acid  (sp.  gr.  1.42)  will  make  5  liters  of  2  N  acid. 
809  cc.  hydrochloric  acid  (sp.  gr.  1.19)  will  make  5  liters  of 

2  N  acid. 

1590  cc.  36%  acetic  acid  will  make  5  liters  of  2  N  acid. 
663  cc.  ammonium,  hydroxide  (sp.  gr.  0.90)  will  make  5  liters 

of  2  N  alkali. 

§  88,  b.  The  solution  of  ammonium  molybdate  is  prepared  as 
follows :  Dilute  200  cc.  of  concentrated  nitric  acid  (sp.  gr.  1.4) 
with  300  cc.  of  water.  Dissolve  75  g.  of  ammonium  molybdate  in 
500  cc.  of  water  and  slowly  pour  this  solution  into  the  dilute 
nitric  acid.  A  precipitate  of  molybdic  acid  will  at  first  precipi- 
tate but  will  afterwards  dissolve. 

§  90.  The  solutions  of  ammonium  sulfocyanate  and  ferric  chlo- 
ride should  be  made  as  follows :  7.5  g.  of  ammonium  sulfocyanate 
dissolved  in  2  liters  of  water ;  6  g.  of  commercial  ferric  chloride 
and  25  cc.  of  hydrochloric  acid  dissolved  in  2  liters  of  water. 

§  92.  Cupromercuric  iodide  (Cu2HgI4).  Treat  a  solution  of 
6.8  g.  of  mercuric  chloride  with  a  solution  of  8.3  g.  of  potassium 
iodide.  Wash  the  precipitate  by  decantation  and  dissolve  it  in  a 
solution  of  8.3  g.  of  potassium  iodide  in  50  cc.  of  water.  Treat 
this  solution  with  a  concentrated  solution  of  12  g.  of  copper  sulfate, 
and  conduct  into  the  solution  a  current  of  sulfur  dioxide  for  the 
reduction  of  the  cupric  salt.  The  deep-red  precipitate  is  cupro- 
mercuric  iodide.  Wash  it  with  water  and  dry  at  a  moderate  tem- 
perature. Yield,  about  20  g.  If  desired,  small  portions  of  this 
preparation  may  be  sealed  in  glass  tubes  and  handed  out  in  this 
condition.  Other  compounds  with  a  change  of  color  at  the  transi- 
tion point  are  silver  mercuric  iodide  (Ag2HgI4),  transition  at 
45°—  50°;  silver  iodide  (Agl),  transition  at  146°;  mercuric  iodide 
(HgI2),  transition  at  127°  but  slow  to  reverse. 

§  115.  Sodium  hypobromite.  A  solution  of  this  compound  suit- 
able for  §  115  may  be  made  according  to  the  following  directions: 
Dissolve  100  g.  of  sodium  hydroxide  in  250  cc.  of  water,  cool,  and 
slowly  add  50  cc.  of  bromine.  Since  the  solution  is  very  corrosive, 
the  directions  on  the  stock  bottle  should  carry  a  caution  to  the 
student. 

[149] 


§  131.  The  most  convenient  thermometer  for  use  in  this  section 
is  the  type  sold  for  use  in  constant-temperature  thermostats.  They 
are  of  small  dimensions,  having  a  scale  reading  from  about  10°  to 
30°,  and  are  graduated  in  tenths  of  a  degree. 

§  213.  The  strips  of  lead  used  in  this  section  should  be  cut 
from  thin  sheet  lead  such  as  can  be  had  in  any  good  tinner's  shop. 
They  should  be  about  1  cm.  wide  and  15  cm.  long. 

NATURAL  GAS  IN  A  BLAST  LAMP 

It  is  hard  to  make  use  of  natural  gas  in  a  blast  lamp,  because 
the  amount  of  air  required  for  combustion  cools  the  flame  below 
the  kindling  temperature,  and  the  flame  is  blown  out.  This  may 
be  remedied  by  diluting  the  gas  with  air  before  it  reaches  the 
burner.  The  easiest  way  to  accomplish  this  is  to  make  a  by- 
pass tube  leading  from  the  hose  supplying  compressed  air  to  the 
one  supplying  gas.  The  connection  can  be  made  by  inserting  a 
T-tube  in  each  hose  and  connecting  the  two  free  ends  with  rubber 
tubing.  A  pinch  clamp  should  be  placed  on  this  tubing  to  regulate 
the  air  admitted.  See  Journal  of  Industrial  and  Engineering 
Chemistry,  VII,  46. 


[150] 


Date  Due 


MAY  20  b31 

APR    23 

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LIBRARY 

COLLEGE   OP   DENTISTRY 
UNIVERSITY   OF  CALIFORNIA 


LIST  OF  THE  ELEMENTS,  THEIR  SYMBOLS, 
AND  ATOMIC  WEIGHTS 


Aluminium  . 
Antimony     . 
Argon 
Arsenic    .     . 

.     Al 
.     Sb 
.     A 
.     As 

27.1 
120.2 
39.9 
74.96 

Molybdenum 
Neodymium  .     . 
Neon     .... 
Nickel  .... 

Mo 
Nd 

Ne 
Ni 

96.0 
144.3 
20.2 
58.68 

Barium     . 

.     Ba 

137.37 

Niton    .... 

Nt 

222.4 

Bismuth  .     . 
Boron 

.     Bi 
.     B 

208.0 
10.9 

Nitrogen   .     .     . 
Osmium    . 

N 
Os 

14.008 
190.9 

Bromine 
Cadmium 

.     Br 
.     Cd 

79.92 
112.4 

Oxygen      .     .     . 
Palladium       .     . 

0 
Pd 

16.00 
106.7 

Caesium    . 
Calcium  .     . 

.     Cs 
.     Ca 

132.81 
40.07 

Phosphorus    . 
Platinum  .     .     . 

P 

Pt 

31.04 
195.2 

Carbon 

.     C 

12.005 

Potassium 

K 

39.10 

Cerium     .     . 
Chlorine  . 

.     Ce 
.     Cl 

140.25 
35.46 

Praseodymium    . 
Radium 

Pr 
Ra 

140.90 
226.00 

Chromium    . 

.     Cr 

52.0 

Rhodium  . 

Rh 

102.9 

Cobalt 

.     Co 

58.97 

Rubidium       .     . 

Rb 

85.45 

Columbium  . 

.     Cb 

93.10 

Ruthenium    . 

Ru 

101.7 

Copper 
Dysprosium  . 
Erbium    .     . 

.     Cu 
.     Dy 
.     Er 

63.57 
162.5 
167.7 

Samarium      .     . 
Scandium  . 
Selenium  .     .     . 

Sa 

Sc 
Se 

150.4 
44.1 
79.2 

Europium 
Fluorine  .     . 
Gadolinium  . 

.     Eu 
.     F 
.     Gd 

152.0 
19.0 
157.3 

Silicon  .... 
Silver    .... 
Sodium      .     .     . 

Si 
Ag 

Na 

28.3 

107.88 
23.00 

Gallium    .     . 

.     Ga 

70.1 

Strontium 

Sr 

87.63 

Germanium  . 

.     Ge 

72.5 

Sulfur  .... 

S 

32.06 

Glucinum 

.     Gl 

9.1 

Tantalum 

Ta 

181.5 

Gold    .     .     . 

.     An 

197.2 

Tellurium       .     . 

Te 

127.5 

Helium     .     . 

.     He 

4.UO 

Terbium    . 

Tb 

159.2 

Holmium 

.     Ho 

163.5 

Thallium  .     .     . 

Tl 

204.0 

Hydrogen     . 
Indium     . 
Iodine 

.     H 

.     In 
.     I 

1.008 
114.8 
126.92 

Thorium    . 
Thulium    .     .     . 
Tin  

Th 
Tm 
Sn 

232.15 
168.5 
118.7 

Iridium    . 

.     Ir 

193.1 

Titanium  . 

Ti 

48.1 

Iron     .     .     . 
Krypton  .     . 
Lanthanum  . 

.     Fe 
.     Kr 
.     La 

55.84 
82.92 
139.0 

Tungsten  .     .     . 
Uranium   . 

Vanadium      .     . 

W 
U 
V 

184.0 
238.2 
51.0 

Lead    .     .     . 

.     PI) 

207.2 

Xenon  .... 

Xe 

130.2 

Lithium  . 

.     Li 

6.94 

Ytterbium      .     . 

Yb 

173.5 

Lutecium 

.     Lu 

175.0 

Yttrium     . 

Yt 

89.3 

Magnesium  . 
Manganese   . 
Mercury  .     . 

.     Mg 
.     Mn 
•     Hg 

24.32 
54.93 
200.6 

Zinc       .... 
Zirconium 

Zn 

Zr 

65.37 
90.6 

